COLLEGE  OF  AGRICULTURE 
DAVIS,  CALIFORNIA 


WORKS  OF  PROF.  F.  A.  GOOCH 

PUBLISHED    BY 

JOHN  WILEY  &  SONS,  Inc. 


Representative  Procedures  in  Quantitative  Chem- 
ical Analysis. 

viii  + 250  pages,    35  figures.      Cloth,  $2.00  net. 

Methods  in  Chemical  Analysis. 

Originated  or  developed  in  the  Kent  Chemical  Lab- 
oratory of  Yale  University.  Compiled  by  Professor 
Frank  Austin  Gooch.  xii  + 536  pages.  6  by  9.  Cloth, 
$4-oo  net.  

BY  GOOCH  AND  BROWNING 

Outlines  of  Qualitative  Chemical  Analysis. 

By  Frank  Austin  Gooch,  Professor  of  Chemistry  in 
Yale  University,  and  Philip  Embury  Browning, 
Assistant  Professor  of  Chemistry  in  Yale  University. 
Third  edition,  revised.  vi  +  i45  pages.  5%  by  8. 
And  colored  section  chart.  Cloth,  $1.25  net. 


Representative  Procedures 

in 

Quantitative  Chemical  Analysis 


BY 


FRANK   AUSTIN   GOOCH 

Professor  of  Chemistry  and  Director  of  the  Kent  Chemical 
Laboratory  in  Yale  University 


FIRST   EDITION 
FIRST    THOUSAND 

UNIVERSITY  OF  CALIFORNIA 

LIBRARY 

COLLEGE  OF  AGRICULTURE 
DAVIS 

- 

NEW  YORK 
JOHN   WILEY    &    SONS,   INC 

LONDON:  CHAPMAN  &  HALL,  LIMITED 
1916 


COPYRIGHT,  1915, 

BY 
FRANK  AUSTIN  GOOCH 


Stanhope  jpress 

F.    H.  GILSON   COMPANY 
BOSTON,  U.S.A. 


PREFACE 


THIS  introduction  to  representative  procedures  in  quanti- 
tative chemical  analysis  is  an  out-growth  of  long  experience 
in  the  teaching  of  that  subject.  Inevitably  it  must  reflect  some 
characteristics  of  many  predecessors,  and  to  all  of  these  the 
author  desires  to  make  every  due  acknowledgment  of  in- 
debtedness. 

Procedures  have  been  discussed  in  relation  to  their  essential 
features,  underlying  principles,  and  varied  applications. 

A  large  amount  of  space  has  been  devoted  to  iodometric 
processes  both  on  account  of  the  beauty  and  accuracy  of  these 
methods,  constantly  growing  in  recognition,  and  because  the 
treatment  in  the  more  elementary  text-books  is  usually  scanty. 

The  arrangement  of  the  subject  matter  is  such  as  to  permit 
the  selection  of  a  narrower  course  or  a  broader  course,  as  may 
be  desired,  without  detriment  to  the  logical  sequence  of  topics. 
Explicit  directions  are  given  for  certain  representative  experi- 
mental processes,  but  if  wider  practice  is  desired  the  general 
description  of  related  processes  should  make  it  possible  to  adapt 
to  these  the  specific  directions  given  for  the  representative  ex- 
perimental processes. 

The  value  of  collateral  reading  in  stimulating  interest  and 
in  amplifying  information  is  recognized,  and,  inasmuch  as 
access  to  the  periodical  literature  is  frequently  difficult,  refer- 
ences in  respect  to  further  information  have  been  made,  so  far 
as  seemed  practicable,  to  a  few  easily  accessible  books,  in  which 
further  reference  to  original  sources  may  usually  be  found,  if 
needed. 


CONTENTS 


CHAPTER  I 

PROCESSES  OF  ANALYSIS 

ANALYTICAL  REACTIONS.  —  The  course  of  chemical  action,  1 ;  successive 
reactions,  3;  non-reversible  and  reversible  reactions,  4;  velocity  of  non- 
reversible  reactions,  5;  completion  of  reversible  reactions,  6;  the  prin- 
ciple of  Le  Chatelier,  9;  empirical  processes,  10. 

CHAPTER  II 
PROCESSES  OF  WEIGHING  AND  MEASURING 

THE  PROCESS  OF  WEIGHING.  —  The  analytical  balance,  11;  determina- 
tion of  the  zero-point,  13;  the  operation  of  weighing,  14;  the  method 
of  equal  swings,  15;  the  method  of  oscillations,  16;  the  sensitiveness  of 
the  balance,  17;  correction  for  inequality  in  the  arms  of  the  balance, 
19;  weighing  by  substitution,  19;  double  weighing,  19;  correction  for 
buoyancy  of  the  air,  19.  The  analytical  weights,  21;  the  testing  of 
weights,  21. 

THE  MEASUREMENT  OF  VOLUME.  —  Volumetric  apparatus,  24;  varia- 
tions of  volume  with  changing  temperature,  25;  correction  of  the  indi- 
cated volume,  26;  calibration  of  instruments,  28;  —  of  flasks,  29;  of 
pipettes,  30;  of  burettes  and  divided  pipettes,  30. 

CHAPTER  III 
PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 

PREPARATION  FOR  WEIGHING.  —  Air  drying,  32;  drying  in  desiccators, 
33;  dehydration  by  heat,  33. 

PROCESSES  DEPENDING  UPON  THE  FORMATION  OF  VOLATILE  PRODUCTS.  — 
Determination  of  water  of  composition,  34;  determination  by  loss  on 
heating,  35;  (experimental  process:  water  in  hydrous  barium  chloride, 
35);  determination  by  loss  on  heating  with  a  flux,  36;  (experimental 
process:  water  in  hydrous  copper  sulphate,  36);  determination  by  direct 
weighing,  36.  Determination  of  carbon  dioxide,  38;  determination  in 

v 


vi  CONTENTS 

carbonates  by  loss  on  heating,  38;  (experimental  process:  carbon  dioxide 
in  calcite,  39);  —  by  loss  on  fusion  with  a  flux,  39;  (experimental  process: 
carbon  dioxide  calcite,  40) ;  by  loss  in  the  action  of  acid  upon  carbonates, 
40;  (experimental  process:  carbon  dioxide  in  calcite,  41);  by  direct  weigh- 
ing, 42.  Determination  of  carbon  as  carbon  dioxide  and  hydrogen  as 
water,  by  the  combustion  process,  44.  Determination  of  boric  acid, 
44;  —  use  of  calcium  oxide  as  the  retainer,  46;  use  of  sodium  tungstate 
as  the  retainer,  47.  Determination  of  the  fixed  product  of  action  of  a 
volatile  reagent,  48;  (experimental  process:  conversion  of  magnesium 
chloride  to  magnesium  sulphate,  48) ;  (experimental  process;  determination 
of  manganese  in  hydrous  manganese  sulphate,  49) . 

ELECTROLYTIC  PROCESSES.  —  The  conditions  of  electrolysis,  49;  stationary 
electrodes,  50;  (experimental  process:  copper  in  copper  sulphate,  51); 
the  rotating  electrode,  52;  (experimental  process:  copper  and  nickel, 
53) ;  small  electrodes  with  weighable  cell,  54. 

PROCESSES  DEPENDENT  UPON  DIFFERENCES  IN  RESPECT  TO  SOLUBILITY.  — 
The  distribution  of  a  solute  between  two  non-miscible  solvents,  56. 
Precipitates  and  conditions  affecting  precipitation,  57;  the  insolubility 
of  precipitates,  58;  nature  of  the  medium,  58;  state  of  subdivision,  58; 
colloidal  suspensions,  60;  purity  of  precipitates,  61;  definiteness  in 
composition  of  precipitates,  62;  the  practical  separation  of  precipitates, 
62;  filtering  and  washing,  62;  filters  of  paper,  63;  —  of  asbestos,  64;  of 
spongy  platinum,  65;  of  anthracene,  66.  Representative  precipita- 
tion processes,  67.  The  precipitation  of  chlorides,  67;  determination 
of  chloride  ion,  67;  (experimental  process:  chlorine  in  barium  chloride, 
69);  (determination  of  bromide  ion  and  iodide  ion,  69);  determination 
of  silver  ion,  69;  (experimental  process:  silver  as  silver  chloride,  70); 
determinations  by  formation  of  other  insoluble  chlorides  —  mercury,  70; 
barium,  71;  aluminium,  72.  The  precipitation  of  sulphates,  73;  de- 
termination of  sulphate  ion,  73;  (experimental  process:  sulphate  ion  in 
copper  sulphate,  76 — filtration  on  paper,  76;  on  asbestos,  77);  general 
procedure,  77;  determination  of  barium,  79;  (experimental  process: 
determination  of  barium  as  sulphate,  80);  determination  of  strontium,  80; 
—  of  calcium,  81;  of  lead,  81;  (experimental  process:  determination  of  lead  as 
sulphate,  81).  The  precipitation  of  phosphates,  81;  determination  of 
orthophosphate  ion,  81;  determination  of  magnesium,  83;  (experimental 
process:  determination  of  magnesium  as  pyrophosphate,  84) ;  determination 
of  manganese,  84;  —  (experimental  process:  determination  of  manganese  as 
pyrophosphate,  86);  determination  of  cadmium  and  zinc,  86.  The  pre- 
cipitation of  oxalates,  86;  precipitation  of  calcium  oxalate,  86;  (experi- 
mental process:  determination  of  calcium  precipitated  as  oxalate,  88); 
precipitation  of  strontium  oxalate,  88;  —  of  barium  oxalate,  88;  of  mer- 
curous  oxalate,  88;  of  copper  oxalate,  89.  The  precipitation  of  carbon- 


CONTENTS  VU 

ates,  89;  precipitation  of  barium,  strontium  and  calcium  carbonates, 
90;  (experimental  process:  determination  of  barium,  90);  precipitation 
of  magnesium-ammonium  carbonate,  90;  (experimental  process:  de- 
termination of  magnesium,  91);  precipitation  of  zinc  carbonate  (basic) 
and  cadmium  carbonate,  91;  precipitation  of  bismuth  basic  "carbonate,- 
92.  The  precipitation  of  hydroxides,  92.  Acidic  hydroxides,  93;  pre- 
cipitation of  silicic  acid  from  silicates,  93;  (experimental  process:  de- 
termination of  silica,  94) ;  determination  of  tungsten  trioxide,  95.  Basic 
hydroxides,  96;  precipitation  of  aluminium  hydroxide,  96;  (experimental 
process:  determination  of  aluminium,  97);  precipitation  of  chromic 
hydroxide,  98;  (experimental  process:  determination  of  chromium,  98); 
precipitation  of  ferric  hydroxide,  99;  (experimental  process:  determina- 
tion of  ferric  hydroxide,  99) ;  the  basic  acetate  process  for  iron  and  alumin- 
ium, 99.  The  precipitation  of  sulphides,  100.  Sulphides  precipitable 
in  acid  solution,  101;  precipitation  of  mercuric  sulphide  by  hydrogen 
sulphide,  101;  (experimental  process:  determination  of  mercury,  102);  pre- 
cipitation of  mercuric  sulphide  from  the  sulphosalt,  102;  precipitation  of 
arsenic  trisulphide,  103;  — of  arsenic  pentasulphide,  103;  of  antimony 
trisulphide,  103;  of  cupric  sulphide,  105;  of  zinc  sulphide,  106;  of 
stannous  sulphide  and  stannic  sulphide,  109;  of  bismuth  sulphide,  109; 
of  lead  sulphide,  109;  of  cadmium  sulphide,  110;  of  molybdenum  sul- 
phide, 110.  Sulphides  precipitable  only  in  alkaline  solution,  110;  pre- 
cipitation of  iron  sulphide,  110;  —  of  manganous  sulphide,  111;  of  nickel 
sulphide  and  cobalt  sulphide,  112. 

CHAPTER  IV 
PROCEDURES;  IN  VOLUMETRIC  ANALYSIS 

NORMAL  SOLUTIONS,  113;  standard  solutions,  114;  varying  normality,  115. 

NEUTRALIZATION  PROCESSES  (Acidimetry  and  Alkalimetry).  —  The  use  of 
indicators,  115;  methyl  orange  (helianthin),  116;  phenolphthalein,  118. 
The  preparation  of  standards,  120;  standard  hydrochloric  acid,  120; 
(experimental  process:  standardization  of  hydrochloric  acid,  121);  stand- 
ard sodium  hydroxide,  121;  (experimental  process:  standardization  of 
sodium  hydroxide,  122);  standard  barium  hydroxide,  124;  (experimental 
processes:  standardization  of  barium  hydroxide,  124;  determination  of 
stronger  acids,  125;  determination  of  sodium  hydroxide  and  sodium  car- 
bonate in  crude  caustic  soda,  126);  determination  of  boric  acid,  127; 
other  standards  and  indicators,  129. 

OXIDATION  PROCESSES.  —  The  use  of  potassium  permanganate,  130;  stand- 
ardization of  potassium  permanganate,  130;  (experimental  process: 
preparation  and  standardization  of  N/W  permanganate,  132);  other 
modes  of  standardizing  permanganate,  133.  Direct  Oxidations  by  the 


viii  CONTENTS 

Permanganate  Solution,  133.  Determination  of  oxalates,  133;  (experi- 
mental process:  determination  of  copper,  133);  determination  of  iron  in 
ferrous  sulphate,  134;  (experimental  process:  determination  of  iron  in 
ferrous  ammonium  sulphate,  135);  determination  of  hydrogen  dioxide, 
135;  determination  of  nitrous  acid  and  nitrites,  136;  determination  of 
manganese  in  manganous  sulphate,  136;  determination  of  formates, 
138;  effect  of  hydrochloric  acid  in  permanganate  titrations,  138.  Oxi- 
dations by  Permanganate  after  a  Preliminary  Reduction,  140.  Reduc- 
tion to  Stable  Condition,  141.  Determination  of  iron  in  ferric  salts 
after  reduction  by  zinc,  141,  by  hydrogen  sulphide,  142,  by  sulphur 
dioxide,  142,  by  stannous  chloride,  143,  by  titanous  sulphate,  143; 
(experimental  process:  iron  in  hematite,  144);  determination  of  vanadic 
acid  reduced  by  sulphur  dioxide,  147.  Reduction  Products  Sensitive 
to  the  Action  of  Air,  137.  Determination  of  molybdic  acid  and  phos- 
phomolybdic  acid,  148;  (experimental  process:  determination  of  phos- 
phorus in  iron  and  steel,  148).  Over-reduction  and  Correction,  150. 
Determination  of  vanadic  acid  reduced  by  zinc,  150;  determination  of 
uranium  in  uranyl  sulphate,  151.  Oxidation  by  Permanganate  of  a 
Reagent  used  in  Excess,  152.  Determination  of  higher  lead  oxides,  152; 
determination  of  higher  oxides  of  manganese,  153;  determination  of 
chlorates,  154;  determination  of  nitrates,  154. 

IODOMETRIC  PROCESSES.  —  Iodine  and  sulphurous  acid,  156;  iodine  and 
thiosulphate,  158;  iodine  and  arsenic  trioxide,  158;  iodine,  iodate 
and  hydrochloric  acid,  160.  The  starch  indicator,  160;  the  end-color, 
160;  the  influence  of  iodides,  161;  preparation  of  the  starch  solution, 
162.  The  preparation  and  use  of  standards,  163;  arsenic  trioxide,  a 
primary  standard,  163;  preparation  of  N/10  arsenite,  163;  the  standard 
iodine  solution,  164;  (experimental  process:  standardization  of  iodine 
against  N/W  arsenite,  164);  the  standard  thiosulphate  solution,  165; 
(experimental  processes:  standardization  of  thiosulphate  against  standard 
iodine,  166;  against  weighed  iodine,  167;  against  iodine  set  free  by 
standard  permanganate,  167);  the  standard  iodate  solution,  168.  The 
determination  of  iodine,  169.  Processes  based  upon  Oxidation  by 
Standard  Iodine,  170.  Direct  titration  by  iodine,  170.  Determination 
of  arsenic  trioxide,  170;  determination  of  thiosulphates,  170;  determi- 
nation of  antimony  trioxide,  171.  Reverse  titration,  171;  determina- 
tion of  sulphurous  acid,  171.  Action  of  iodine  in  excess  and  titration  of 
the  excess,  172;  determination  of  soluble  sulphides,  172.  Processes 
based  upon  Liberation  of  Iodine  in  Solution,  173.  Determination  of 
bromine,  173;  determination  of  chlorine,  174;  the  reaction  between  an 
iodide  and  an  iodate  in  presence  of  an  acid,  174;  determination  of 
iodates,  175;  (experimental  processes:  determination  of  iodate  by  thio- 
sulphate titration  of  liberated  iodine,  175;  by  arsenite  absorption  of 
iodine,  175);  determination  of  periodates,  176;  determination  of  iodates 


CONTENTS  IX 

and  periodates  in  association,  177;  determination  of  iodides,  177;  action 
of  iodate  in  presence  of  strong  hydrochloric  acid,  178;  (experimental 
processes:  determination  of  iodide  from  iodide  set  free,  179;  — from  excess 
of  iodate,  179;  by  direct  action  of  standard  iodate,  180).  Other  substances 
(oxidizers)  reactive  with  iodide  or  hydriodic  acid,  180;  determination 
of  hydrogen  dioxide,  181;  —  of  hypochlorites,  181;  of  bromates,  181;  of 
chromates,  182;  of  lead  dioxide,  182;  of  cerium  dioxide,  183;  of  gold 
trichloride,  183;  of  ferric  chloride,  184;  of  copper  in  cupric  salts,  184. 
Titration  of  Iodine  after  Extraction  from  Aqueous  Solution,  185.  De- 
composition of  iodides  by  nitrous  acid,  186.  Processes  depending  upon 
the  Determination  of  Iodine  Volatilized  to  the  Distillate,  186.  Deter- 
mination of  iodine  in  iodides,  187;  approximate  determination  of  arsenic, 
antimony,  vanadium,  molybdenum,  cerium,  and  selenium,  188.  The 
Volatilization  of  Iodine  in  Acid  Solution  and  the  Reoxidation  of  the 
Residue  by  Iodine  in  Alkaline  Solution,  189.  Determination  of  iodides 
by  the  action  of  arsenic  acid  and  reoxidation  by  iodine,  189;  (experi- 
mental process:  determination  of  iodine  in  an  iodide,  190);  determina- 
tion of  arsenic  in  arsenic  acid,  190;  (experimental  processes:  determina- 
tion of  arsenic  in  arsenic  acid,  191);  determination  of  antimony  in  anti- 
monic  acid,  192.  Evolution  of  Chlorine  and  Absorption  by  Potassium 
Iodide:  Titration  of  the  Iodine  set  Free,  193;  the  operation,  193. 
Determination  of  manganese  dioxide,  194;  —  of  chlorates,  194;  of 
nitrates,  194;  of  nitrites,  196.  Evolution  of  Bromine  and  Absorption 
by  Potassium  Iodide,  196;  Titration  of  the  Iodine  set  Free,  196.  De- 
termination of  vanadates,  196;  —  of  selenates,  197;  of  tellurates,  198. 
lodometric  Processes  based  upon  a  Preliminary  Reduction  or  Oxidation, 
198.  Determination  of  vanadates,  reduced  by  oxalic,  tartaric,  or  citric 
acid,  198;  determination  of  bromates,  by  reduction  with  arsenious  acid 
and  titration  of  the  excess,  199;  determination  of  chromates,  by  reduc- 
tion with  arsenious  acid  and  titration  of  the  excess,  200;  determination 
of  tellurous  oxide,  by  action  of  permanganate  and  iodide  with  titration 
of  iodine  set  free,  201.  Processes  based  upon  the  Reaction  between 
Acids  and  an  lodide-Iodate  Mixture,  202.  Determination  of  hydro- 
chloric acid  and  sulphuric  acid,  202;  determination  of  boric  acid,  202; 
(experimental  process:  determination  of  boric  acid,  203);  determination 
of  acids  liberated  in  the  hydrolysis  of  salts,  203;  —  of  aluminium  sul- 
phate and  aluminium  chloride,  204;  of  other  salts  showing  similar 
behavior,  204;  determination  of  alkali  hydroxides,  204;  determination 
of  alkali  carbonates,  205.  Processes  dependent  upon  the  Action  of 
Iodine  upon  Basic  Hydroxides  to  form  an  Iodide  and  an  Iodate  in 
Mixture,  205.  Determination  of  basic  hydroxides,  205.  Differential 
determination  of  acids,  207;  —  of  hydrochloric  acid,  207;  of  sulphuric 
acid,  208;  of  carbonic  acid,  208.  Processes  based  upon  the  Action  of 
Potassium  Iodate  in  Hydrochloric  Acid  upon  Substances  Susceptible  to 


X  CONTENTS 

Oxidation,    208;    (experimental  processes;    determination  of  sulphurous 
add,  209;  determination  of  copper  in  ore,  209). 

PRECIPITATION  PROCESSES.  —  The  End-point  marked  by  Absence  of 
Further  Precipitation,  210.  Determination  of  silver  by  precipitation 
as  silver  chloride,  210;  determination  of  tellurous  acid  by  precipitation 
of  tellurous  iodide,  211.  The  End-point  marked  by  the  Beginning  of 
Precipitation,  212.  Determination  of  the  cyanide  ion,  212.  The  End- 
point  marked  by  a  New  Reaction,  213.  Determination  of  the  chloride 
ion  by  precipitation  as  silver  chloride,  with  an  indicator  in  the  solution, 
213;  determination  of  silver  by  precipitation  as  silver  thiocyanate, 
with  a  ferric  indicator  in  the  solution,  214;  (experimental  process:  pre- 
cipitation of  silver  as  silver  thiocyanate,  214).  The  End-point  marked  by 
Drop  Tests  outside  the  Solution,  215.  Determination  of  phosphoric 
acid  by  precipitation  as  uranyl  phosphate  with  a  ferro-cyanide  indicator 
outside  the  solution,  215. 

GASOMETRIC  PROCESSES.  —  Gas- Volumetric  Methods,  217.  Determina- 
tion of  nitrates,  217;  determination  of  nitrogen  in  ammonium  salts,  219; 
determination  of  nitrogen  in  organic  substances.  Methods  of  Gas- 
Analysis,  220.  Determination  of  oxygen  and  of  nitrogen  (with  the  inert 
elements)  in  air,  221. 

COLORIMETRIC  PROCESSES.  —  Comparison  tubes,  222;  colorimeters,  223; 
the  nephelometer,  224. 

CHAPTER  V 
SYSTEMATIC  ANALYSIS 

SEPARATIONS  AND  DETERMINATIONS.  —  The  preparation  of  the  substance, 
226;  qualitative  tests,  227;  choice  of  methods,  227;  grouping  of  basic 
ions,  228;  typical  forms  precipitated  by  group  reagents,  229.  Repre- 
sentative Methods  of  Separation.  —  Analysis  of  brass,  230;  (experimen- 
tal process:  determination  of  copper,  zinc,  tin,  lead,  iron,  and  nickel  in 
brass;  statement  of  analysis,  230).  Analysis  of  limestone,  233;  (experi- 
mental process:  analysis  of  limestone;  determination  of  loss  on  ignition, 
carbon  dioxide,  ferric  oxide,  ferrous  oxide,  insoluble  siliceous  matter,  silica, 
alumina  (etc.),  calcium  oxide,  and  magnesium  oxide;  statement  of  analysis, 
233).  Analysis  of  silicates,  237;  determination  of  alkali  elements,  238; 
(experimental  process:  determination  of  potassium  and  sodium  in  feldspar, 
239);  determination  of  fluorine,  240.  Analysis  of  substances  yielding 
ammonia,  242. 

INDIRECT  ANALYSIS.  —  Analysis  of  chlorides  and  bromides,  243;  (ex- 
perimental process:  indirect  determination  of  chlorine  and  bromine,  244). 
Analysis  of  nitrates  and  nitrites,  245. 

INTERNATIONAL  ATOMIC  WEIGHTS,  1914. 


REPRESENTATIVE  PROCEDURES 

IN 
QUANTITATIVE   CHEMICAL   ANALYSIS 


CHAPTER  I 
PROCESSES  OF  ANALYSIS 

THE  methods  of  Quantitative  Chemical  Analysis  are  generally 
based  upon  definite  chemical  transformations.  In  Gravimetric 
Processes  products  of  reaction  are  isolated  and  weighed;  in 
Volumetric  Processes  transformations  are  brought  about  by 
means  of  measured  solutions  of  known  content  by  weight.  A 
few  processes  of  quantitative  determination  depend  upon  the 
measurement  of  certain  physical  characteristics,  such  as  the 
volume  filled  under  determined  conditions  of  temperature  and 
pressure  by  gaseous  substances  or  the  effect  of  light  upon  solu- 
tions —  color,  opalescence,  refractive  power,  polarization. 

Analytical  Reactions 

Not  all  the  reactions  which  are  useful  in  identifying  individual 
elements  (elementary  analysis)  or  compounds  (proximate  analy- 
sis) are  adapted  to  quantitative  determinations.  Reactions 
must  be  so  chosen,  and  so  utilized,  that  substances  may  either 
be  separated  in  definite  condition  for  weighing  or  measurement, 
or  else  may  be  made  to  react  in  well-determined  ways  with 
known  amounts  of  the  reagents.  In  choosing  and  applying  re- 
actions for  the  purposes  of  quantitative  analysis  it  is,  of  course, 
important  to  keep  in  mind  the  fundamental  and  general  prin- 
ciples of  chemical  reaction. 

The  Course  of  Chemical  Action.  —  When  chemical  systems 
undergo  change  the  tendency  to  reaction  is  directed  in  lines 

1 


2  QUANTITATIVE  CHEMICAL  ANALYSIS 

which  correspond  to  the  development  of  free  energy  —  that  is, 
energy  available  for  the  performance  of  work.  The  changes  of 
energy  which  accompany  chemical  reaction  may  involve  internal 
changes  in  the  system  as  well  as  the  evolution  of  free  energy. 
As  heat  becomes  latent  in  the  change  from  ice  to  liquid  water, 
or  of  liquid  water  to  water  vapor,  so  some  of  the  chemical  energy 
of  reaction  may  disappear  in  the  internal  rearrangement  of  the 
system.  But  it  appears  that  all  processes  which  proceed 
spontaneously  after  having  been  started  —  that  is,  processes 
which  proceed  without  continual  supplying  of  energy  from 
without  the  system  —  are  accompanied  by  the  evolution  of  free 
energy,  the  least  stable  condition  of  a  system  undergoing 
spontaneous  reaction  being  that  in  which  the  system  is  capable 
of  evolving  the  greatest  amount  of  free  energy,  and  the  most 
stable  condition  that  which  corresponds  to  the  greatest  eVolu- 
tion  of  free  energy.  If  it  be  assumed,  as  is  probable,  that  the 
free  energy  and  the  total  heat  of  reaction  have  the  same  value 
in  the  neighborhood  of  the  absolute  zero  of  temperature,  then, 
all  reactions,  once  started,  should  proceed  at  such  temperatures 
with  evolution  of  heat;  and,  at  such  temperatures,  the  amount 
of  heat  evolved  should  measure  the  resultant  of  the  chemical 
affinities,  or  the  tendency  to  reaction.  The  interval  between 
the  ordinary  atmospheric  temperature  and  the  absolute  zero  is, 
however,  not  very  great  in  comparison  with  the  entire  available 
range  of  reaction  temperatures,  and  so  it  is  to  be  expected  that 
reactions  at  ordinary  temperatures  will  generally  proceed  in  a  way 
to  produce  the  maximum  evolution  of  heat.  As  a  matter  of  fact, 
heats  of  reaction  may  often  be  used  to  indicate  the  direction  in 
which  a  reaction  involving  certain  elements  or  ions  will  move. 
For  example,  amounts  of  heat  evolved  in  the  neutralization 
of  a  gram-equivalent  of  hydrochloric  acid  and  of  a  gram- 
equivalent  of  acetic  acid  by  the  action  of  sodium  hydroxide  are, 
respectively,  13,700  calories  and  13,400  calories.  It  is,  there- 
fore, natural  to  suppose  that  upon  gradually  adding  the  water 
solution  of  a  gram-equivalent  of  sodium  hydroxide  to  the  water 
solution  containing  a  gram-equivalent  of  hydrochloric  acid  and 
a  gram-equivalent  of  acetic  acid  the  hydrochloric  acid  will  be 


PROCESSES  OF  ANALYSIS  3 

neutralized  before  the  acetic  acid,  and  this  expectation  is  very 
closely  realized.  The  residue  left  after  the  evaporation  of  such 
a  solution  consists  almost  entirely  of  sodium  chloride. 

Similarly,  when  a  mixture  of  two  gram-molecules  of  hydrogen 
and  one  gram-molecule  of  oxygen  is  acted  upon  at  the  ordinary- 
temperature  by  enough  outside  energy  to  start  the  combination 
of  these  elements,  the  reaction 

2  H2  +  02  =  2  H20 

(which  evolves  68,360  calories,  assuming  that  the  water  formed 
is  allowed  to  condense  to  the  liquid  phase)  proceeds  to  comple- 
tion with  evolution  of  heat  unless  the  resulting  temperature  of 
the  mass  be  cooled  artificially  below  a  definite  point  (about  560 
degrees)  at  which  the  chemical  change  is  spontaneous. 

On  the  other  hand,  reactions  may  proceed  in  the  direction 
which  implies  absorption  of  heat  when  heat  is  supplied  at 
sufficiently  high  temperatures  from  without  the  system;  and 
at  intermediate  temperatures  an  equilibrium  between  opposite 
tendencies  may  be  established.  For  example,  the  reaction  in 
which  hydrogen  and  oxygen  form  water  with  evolution  of  heat 
at  560  degrees, 

2H2  +  02^2H20, 

shows  a  reversal  at  temperatures  of  about  1000  degrees,  and  it 
proceeds  equally  in  opposite  directions  at  about  2500  degrees. 
Similarly,  the  reaction  between  a  gram-equivalent  of  hydro- 
chloric acid  and  a  gram-equivalent  of  sodium  hydroxide,  which 
results  at  ordinary  temperatures  in  practically  complete  mutual 
neutralization,  is  markedly  reversible  at  temperatures  approach- 
ing low  red  heat : 

NaOH  +  HC1  <=*  NaCl  +  H20. 

Successive  Reactions.  —  When  a  system  is  undergoing 
transformation,  reactions  yielding  systems  of  successively 
greater  degrees  of  stability  may  intervene  between  the  initial 
condition  and  the  final  condition.  For  example,  the  reactions 
in  a  system  made  up  of  barium  dioxide,  potassium  dichromate, 
sulphuric  acid  and  water  may  take  place  in  successive  stages. 


4  QUANTITATIVE  CHEMICAL  ANALYSIS 

First,  there  is  the  production  of  hydrogen  dioxide  and  chromic 
acid  in  presence  of  sulphuric  acid;  secondly,  the  formation  of 
perchromic  acid;  and,  thirdly,  the  reduction  of  the  per  chromic 
acid  with  formation  of  chromic  sulphate. 

Ba02  +  K2Cr207  +  6  H2SO4 

=  BaSO4 .+  2  KHSO4  +  3  H2SO4  +  H2Cr2O7  +  H202 
=  BaSO4  +  2  KHS04  +  3  H2SO4  +  H2Cr2O8  +  H2O 
=  BaS04  +  2  KHS04  +  Cr2(S04)3  +  5  H20  +  2  O2. 

In  exact  quantitative  analysis  it  is  important  to  make  use  of 
reactions  and  conditions  which  result  in  the  establishment  of 
definite  and  stable  systems  at  once,  or  at  least  within  a  reasonable 
time. 

Non-Reversible  and  Reversible  Reactions.  —  Processes 
which  result  in  the  formation  of  inactive  products  and  which  are 
non-reversible  proceed  to  completion,  when  once  started,  with- 
out regard  to  the  presence  of  reaction  products.  On  the  other 
hand,  processes  in  which  all  products  are  active,  and  which  are 
therefore  reversible,  are  affected  not  only  by  adjustment  of 
conditions  or  by  the  increase  of  a  factor  but  by  the  removal  of 
an  active  product. 

Examples  of  processes  which  rest  upon  non-reversible  reac- 
tions and  which  proceed  to  completion  uninfluenced  by  the 
amounts  of  the  products  of  reaction  are  the  ignition  of  silicon, 
aluminium,  and  iron  hydroxides,  and  of  ammonium  magnesium 
phosphate, 

Si02  •  nH20  =  Si02  +  nH20, 
2  (AIO)OH       =  A1203  +  H20, 
2  Fe03H3          =  Fe2O3  +  3  H20, 
2  NH4MgP04  =  Mg2P207  +  2  NH3  +  H20; 

the  bleaching  of  iodine  by  sodium  thiosulphate, 

I2  +  2  Na2S203  =  2  Nal  +  Na2S2O6; 

the  precipitation  of  manganic  hydroxide  in  the  reaction  between 
manganese  sulphate,  and  ammonium  hydroxide,  and  bromine, 


PROCESSES  OF  ANALYSIS  5 

2  MnS04  +  8  NH4OH  +  2  Br2  =  H2Mn2O5  +  4  NKUBr 

+  2(NH4)2SO4  +  3H2O; 

and  the  reduction  of  ferric  chloride  by  oxalic  acid, 

2  FeCl3+  H2C204  =  2  FeCl2  +  2  HC1  +  2  CO2. 

Velocity  of  Non-Reversible  Reactions.  —  The  velocity 
with  which  a  non-reversible  reaction  reaches  its  end  depends 
upon  the  masses  of  the  reacting  substances  and  the  affinity 
factors  involved.  The  rate  of  a  non-reversible  reaction,  when 
slow  enough  to  be  measurable,  is  generally  accelerated  by  elevation 
of  temperature,  and  at  each  instant  it  is  found  to  be  proportional 
to  the  concentration  of  the  reacting  substances.  For  example,  the 
reaction  between  iodine  and  sodium  thiosulphate,  mentioned 
above,  is  practically  immediate  at  the  ordinary  atmospheric 
temperature;  that  between  ferric  chloride  and  oxalic  acid  at 
a  similar  temperature  and  in  dilute  solution  may  require  years 
for  its  completion,  but  this  reaction  may  be  pushed  to  comple- 
tion in  a  few  minutes  by  raising  the  temperature  to  100  degrees. 
The  proper  conditions  in  respect  to  the  times  and  the  tempera- 
tures required  to  effect  as  completely  as  possible  the  transforma- 
tion of  a  system  must  be  found  experimentally  in  each  individual 
case. 

As  the  non-reversible  process  advances  the  velocity  of  reaction 
diminishes,  and  sometimes  to  such  an  extent  that  it  is  not 
feasible,  in  an  analytical  process,  to  continue  the  reaction  to 
completion.  In  the  case  of  such  a  slow  non-reversible  reaction 
it  is  sometimes  possible  to  push  the  velocity  of  reaction  by  adding 
a  large  excess  of  the  active  reagent  and  then  to  determine  by 
some  suitable  process  the  amount  of  that  excess.  For  example, 
in  the  reaction  between  potassium  permanganate  and  selenious 
acid  (in  solution  with  dilute  sulphuric  acid)  the  rate  of  the 
reaction 

5  H2SeO3  +  2  KMnO4  +  3  H2SO4  =  5  H2SeO4  +  K2S04 
+  2  MnSO4  +  3  H20 

is  exceedingly  slow  and  the  straight  course  of  the  reaction  is 
complicated  by  the  intervention  of  a  secondary  reaction  in 


6  QUANTITATIVE  CHEMICAL  ANALYSIS 

which  a  portion  of  the  permanganate  acts  upon  the  manganous 
salt  with  formation  of  higher  oxides  of  manganese.  By  adding 
a  large  excess  of  the  permanganate  the  selenious  acid  may  be 
more  rapidly  oxidized,  and  that  excess  together  with  any  highei 
oxides  of  manganese  may  then  be  reduced  by  a  known  amounl 
of  oxalic  acid,  leaving  an  excess  of  the  last  to  be  destroyed  by 
the  direct  action  of  more  of  the  permanganate. 

The  Completion  of  Reversible  Reactions.  —  In  reversibh 
reactions,  the  result  is  influenced  by  the  concentrations  of  iht 
products  as  well  as  by  the  concentrations  of  the  factors,  in  accord- 
ance with  the  LAW  OF  MASS-ACTION. 

When  two  substances  A  and  B  react  to  form  two  other  sub- 
stances C  and  D,  the  resulting  concentrations  of  the  substances 
expressed  in  terms  of  gram-molecules  may  be  represented  by 
the  terms  CA,  CB,  Cc  and  CD.  When  the  reaction  takes  place 
in  single  molecular  proportions  (mols)  the  forward  action  anc 
the  reverse  action  may  be  represented  by  the  expressior 
k  •  CA  •  CB  and  ki  •  Cc  •  CD,  respectively,  in  which  k  and  k\  stanc 
for  affinity  constants,  and  the  equation  of  equilibrium  is 


When  the  substances  A  and  B  react  in  multiple  molecular  pro- 
portions the  numbers  representing  molecules  ni,  n%,  n%  and  n± 
respectively,  appear  as  exponents,  and  the  equation  takes  the 
form 

CA^-CB*      fci       „ 

-  —    -  •   —     /\ 

CV'-CV*      k 

When  the  molecular  units  of  substances  are  subject  to  change 
apart  from  the  chemical  reaction  (as  by  dissociation,  hydrolysis 
polymerization,  or  depolymerization  in  solution)  so  that  the 
numbers  of  active  units  are  changed,  such  changes  affect  the 
conditions  of  equilibrium.  Representing  the  coefficients  ol 
unitary  change  (as  determined  by  osmotic,  cryoscopic,  ebullio- 
scopic,  or  conductivity  measurements)  as  ii,  i%,  is,  and  i4,  the 
equation  of  equilibrium,  at  constant  temperature  and  volume 
takes  for  the  substances  reacting  in  single  molecular  proportions 
the  form 


PROCESSES  OF  ANALYSIS  7 

and  for  cases  of  reaction  in  multiple  molecular  proportions, 
represented  by  n\  and  n2,  etc.,  the  expression  takes  the  form 


Ccn^  •  CV*      k 

The  validity  of  these  expressions  has  been  shown  experimen- 
tally in  a  number  of  cases.  For  example,  it  has  been  found 
that  when  solutions  of  sodium  sulphate  and  nitric  acid  are 
brought  together  in  a  calorimeter  the  reaction 

Na2S04  +  HN03  <=*  NaHSO4  +  NaN03, 

when  interpreted  in  accordance  with  the  heat  effects,  proceeds 
in  accordance  with  the  expression 


. 

D  NaHSO4  *      NaNO3 

where  the  numerical  exponents  represent  the   coefficients  of 
unitary  change  of  the  individual  substances  in  solution. 

Since  the  concentrations  of  insoluble  substances  may  be 
regarded  as  practically  constant,  terms  relating  to  them  may 
be  eliminated  from  the  equation  of  equilibrium.  For  example, 
in  the  reaction  which  takes  place  between  dissolved  carbon 
dioxide  and  insoluble  barium  carbonate, 


BaC03  +  (C02  +  nH20)  ^BaH2(C03)2  +  (n  -  1)H20, 

the  amount  of  the  solvent  (water)  being  large  and  practically 
constant,  the  equation  of  equilibrium  may  be  written 


and  this  is  found  to  be  in  accord  with  the  results  of  experiment. 
Similarly,  in  the  reaction  between  insoluble  zinc  sulphide  and 
sulphuric  acid  with  the  production  of  zinc  sulphate  and  hydro- 
gen sulphide, 

ZnS  +  H2SO4  <=»  ZnSO4  +  H2S, 

the  value  of  i  for  the  soluble  substances  being  respectively  2.06, 
0.98,  and  1.04,  under  the  conditions  of  action,  the  equation  is 

^H2S04  ^ 

^0.98        .  r/1.04    ~~  *^> 


8  QUANTITATIVE  CHEMICAL  ANALYSIS 

and  since  the  zinc  sulphate  and  hydrogen  sulphide  are  always 
produced  in  equal  molecular  concentrations,  the  expression 
takes  the  form 


^H2SO4 
2.02 


which  agrees  with  the  experimental  results.* 

In  practical  analysis,  exact  knowledge  is  often,  and  indeed 
generally,  wanting  in  respect  to  the  data  necessary  for  the  ex- 
pression of  quantitative  effects  under  given  conditions;  but 
the  general  form  of  the  equation  will  indicate  qualitatively  the 
directions  in  which  changes  of  concentration  take  effect. 

For  example,  in  the  reaction  between  barium  chloride  and 
sulphuric  acid, 

BaCl2  +  H2S04  <=*  BaSO4  +  2  HC1, 

it  is  obvious  that  an  increase  in  the  concentration  of  sulphuric 
acid  will  tend  to  make  the  reaction  complete  in  respect  to  the 
removal  of  barium  chloride,  and  that  an  increase  in  the  con- 
centration of  barium  chloride  will  tend  to  complete  the  reaction 
in  respect  to  the  removal  of  sulphuric  acid,  and  that  an  increase 
in  the  concentration  of  hydrochloric  acid  will  tend  to  render 
each  of  these  reactions  less  complete.  Experience  shows 
that  the  concentration  of  the  barium  sulphate  in  solution  is 
very  small  (0.0023  grm.  in  1  liter  of  water),  remaining  prac- 
tically constant  and  being  only  slightly  affected  by  small 
changes  in  the  concentration  of  the  hydrochloric  acid.  More- 
over, it  is  found  that  the  solvent  effect  of  very  large  concentra- 
tions of  hydrochloric  acid  may  be  balanced  by  comparatively 
small  concentrations  of  sulphuric  acid,  which  means,  of  course, 
that  the  tendency  toward  the  reverse  action  (2  HC1  :  BaSOJ  is 
very  small  in  comparison  with  the  tendency  toward  the  forward 
action  (BaCl2  :  H2S04). 

A  similar  case  is  that  involving  the  precipitation  of  silver 
chloride  by  the  reaction  between  sodium  chloride  and  silver 
nitrate  for  the  purpose  of  determining  either  the  silver  of  the 

*  Chesneau:  Theoretical  Principles,  trans.  Lincoln  and  Carnahan,  p.  79. 


PROCESSES  OF  ANALYSIS  9 

silver  nitrate  or  the  chlorine  of  the  sodium  chloride.     These 
processes  are  based  upon  the  reversible  reaction 
AgNO3  +  NaCl  *±  AgCl  +  NaN03. 

The  concentration  of  the  nearly  insoluble  silver  chloride  is  ver-y 
small  and  practically  constant,  but  a  slight  increase  in  the 
amount  of  either  sodium  chloride  or  silver  nitrate  will  determine 
further  precipitation  of  silver  chloride.* 

The  Principle  of  Le  Chatelier.  —  In  respect  to  the  effect  which 
may  be  induced  by  variations  of  temperature  or  pressure  or  con- 
centration in  reversible  reactions,  the  Principle  of  Le  Chatelier 
affords  a  useful  criterion.  According  to  this  generalization  any 
change  made  from  the  outside  in  the  factors  of  equilibrium  of  a 
balanced  system  is  followed  by  a  change  of  the  opposite  sort  in 
the  similar  factors  of  equilibrium  within  the  system.  Thus, 
an  increase  in  the  external  temperature  will  favor  a  reaction  in 
which  heat  disappears  within  the  system  and  will  oppose  an 
action  in  which  heat  is  evolved;  while  a  lowering  of  the  tem- 
perature will  tend  to  turn  the  reaction  in  the  direction  which 
implies  evolution  of  heat.  An  increase  of  pressure  upon  the 
system  will  favor  the  course  of  reaction  in  which  the  pressure 
(or  volume  under  constant  pressure)  of  the  system  will  be 
lessened;  while  diminution  of  pressure  upon  the  system  will  aid 
the  course  of  reaction  in  which  the  pressure  (or  volume  under 
constant  pressure)  of  the  system  tends  to  increase. 

For  example,  in  an  equilibrium  between  calcium  carbonate, 
calcium  oxide,  and  carbon  dioxide  in  the  reversible  reaction 

With  absorption  of  heat 


CaCO3         4 CaO  +  CO2 

With  evolution  of  heat 

the  application  of  heat  will  tend  to  break  up  the  carbonate, 
while  cooling  will  favor  the  reaction  by  which  the  carbonate  is 
formed  from  the  oxides.  In  this  reaction,  in  which  the  heat 
effect  is  considerable,  the  attainment  of  equilibrium  is  rapid  and 
markedly  dependent  upon  the  temperature;  but  when  the  heat 
of  a  reaction  is  small  the  influence  of  heating  or  cooling  is  not 
*  See  pp.  67,  69. 


10  QUANTITATIVE  CHEMICAL  ANALYSIS 

immediate  or  marked,  as  in  the  esterification  of  alcohol  with 
acetic  acid, 

C2H5OH  +  HOC2H30  <±  C2H5OC2H30  +  H20, 

the  final  equilibrium  of  this  reaction  being  attained  at  the  at- 
mospheric temperature  only  after  several  days,  and  at  100  de- 
grees only  after  some  hours. 

Similarly,  in  the  adjustment  of  equilibrium,  under  changing 
external  pressure  an  increase  of  pressure  upon  the  system  will 
favor  that  process  which  implies  diminution  of  pressure  (or  vol- 
ume) of  the  system  while  relief  of  pressure  upon  the  system  will 
tend  to  promote  the  action  by  which  greater  pressure  (or  vol- 
ume) of  the  system  will  be  produced.  For  example,  in  the  equi- 
librium between  calcium  carbonate,  calcium  oxide,  and  carbon 
dioxide  increase  of  pressure  promotes  the  formation  of  the  solid 
carbonate  while  relief  of  pressure  favors  the  decomposition  of 
the  carbonate  into  calcium  oxide  and  gaseous  carbon  dioxide. 

With  diminished  pressure 

CaCOs          ]          *          CaO  +  CO2 

With  increased  pressure 

The  reaction  between  alcohol  and  acetic  acid,  which  involves  no 
significant  change  in  volume,  is  independent  of  the  pressure. 

Furthermore,  as  has  been  illustrated  in  the  previous  dis- 
cussion of  mass-action  in  relation  to  reversibility,  an  increase 
in  the  concentration  of  one  component  favors  the  formation 
of  another  component  (or  phase)  involving  a  decrease  in  the 
concentration  of  the  first,  and  a  decrease  in  the  concentration 
of  a  component  favors  the  formation  of  another  component 
(or  phase)  involving  an  increase  in  the  concentration  of  the  first. 

Empirical  Processes.  —  Sometimes  it  is  a  matter  of  con- 
venience to  make  use  of  a  process  for  which  it  has  been  estab- 
lished that  correct  results  may  be  indicated  under  a  definite 
set  of  empirical  conditions  rather  than  a  process  which  depends 
upon  the  completion  of  a  definite  reaction  in  accordance  with 
underlying  theory;  but  such  processes  are  employed  only  when 
speed  rather  than  the  highest  accuracy  is  the  object  in  view  or 
as  a  last  resort  in  operations  demanding  exactness. 


CHAPTER  II 
WEIGHING  AND  MEASURING 

The  fixing  of  the  amount  of  a  substance  taken  for  analysis,  or 
found  in  the  course  of  analysis,  depends  either  directly  upon  an 
exact  process  of  weighing  or,  in  the  case  of  a  liquid  or  gas  for  which 
the  relation  of  weight  to  volume  is  known,  upon  the  measurement 
of  volume.  Excepting  processes  of  gas-analysis,  in  which  the 
proportionate  composition  of  a  substance  is  often  stated  in  vol- 
umes, the  results  of  analysis  are  ultimately  stated  in  terms  of 
weight,  and  generally  in  percentages  by  weight. 

The  Process  of  Weighing 

The  Analytical  Balance.  —  The  analytical  balance  (Fig.  1), 
by  which  weights  are  determined,  has  as  essential  parts  a  beam 
which  is  supported  at  its  middle  point  by  a  knife-edge  set  perpen- 
dicularly and  resting  upon  a  plane  bearing  at  the  top  of  a  support- 
ing pillar,  a  pan  suspended  at  each  end  of  the  beam  from  an 
upturned  knife-edge  fixed  in  the  beam  and  parallel  to  the  middle 
knife-edge,  and  a  pointer  reaching  perpendicularly  from  the 
middle  knife-edge  to  a  divided  scale  so  as  to  show  the  movements 
of  the  beam.  A  case  with  glass  sides  protects  the  balance  from 
dust  and  disturbing  currents  of  air  during  the  process  of  weigh- 
ing, and  a  beam-support  and  pan-supports  which  may  be  set  and 
released  by  suitable  mechanism  are  important  adjuncts.  Con- 
stancy and  sensitiveness  are  characteristics  necessary  to  good 
service  on  the  part  of  the  balance;  and  these  essentials  are 
secured  by  proper  attention  to  details  of  construction  and 
adjustment,  such  as  rigidity  and  lightness  of  the  beam,  the 
preservation  of  a  proper  relation  between  the  center  of  gravity 
of  the  loaded  beam  and  the  plane  of  the  supporting  knife-edge, 
and  the  proper  placing  of  the  knife-edges  in  respect  to  one 
another. 

11 


12 


QUANTITATIVE  CHEMICAL  ANALYSIS 


The  beam  must  be  rigid  enough  to  keep  its  shape  under  the 
stress  of  all  reasonable  loading  and  light  enough  to  respond  to 
slight  variations  in  the  load.  The  center  of  gravity  of  the  beam 
with  pans  attached  should  fall  below  the  horizontal  plane  in 
which  the  supporting  knife-edge  lies,  practically  the  plane  of 


FIG.  1. 

the  bearing;  for  otherwise  the  balance  will  be  unreliable,  being 
unstable  when  the  center  of  gravity  lies  above  the  plane  of  the 
knife-edge,  and  indifferent  when  the  center  of  gravity  lies  in  the 
plane  of  the  knife-edge.  But  the  distance  between  the  plane 
of  the  knife-edge  and  the  center  of  gravity  should  be  very  small, 
else  small  variations  in  the  load  will  be  without  appreciable 
effect;  while  it  must  not  be  too  small  lest  the  oscillations  of  the 
beams  be  tediously  slow.  In  a  reasonably  sensitive  instrument 
the  distance  between  the  center  of  gravity  of  the  beam  and  the 
plane  of  the  knife-edge  probably  approximates  0.0025  mm. 

The  terminal  knife-edges  for  pan  suspension  and  the  middle 
knife-edge  for  the  beam-suspension,  all  parallel  to  one  another 
and  perpendicular  to  the  beam,  should  lie  in  the  same  plane; 
for  otherwise,  with  the  beam  rigid,  the  relation  between  the 
position  of  the  center  of  gravity  and  the  plane  of  the  beam 
knife-edge  will  be  subject  to  change  with  variation  of  the  load 


WEIGHING  AND  MEASURING  13 

upon  the  pans.  The  balance  will  lose  sensitiveness  as  the  load 
increases,  the  center  of  gravity  moving  downward,  in  case  the 
plane  of  the  pan  knife-edges  is  lower  than  the  plane  of  the  middle 
knife-edge;  it  will  become  inert  and  then  unstable  as  the  loacL 
increases  and  moves  the  center  of  gravity  upward  in  case  the 
plane  of  the  pan  knife-edges  is  above  the  plane  of  the  middle 
knife-edge.  It  is  to  be  noted,  however,  that  lack  of  perfect 
rigidity  of  beam  and  imperfect  adjustment  of  the  knife-edges  to 
the  same  plane  may  to  some  extent  be  mutually  corrective,  and 
that  varying  sensitiveness,  due  to  flexure  of  the  beam  under  an 
increased  load,  may  be  counteracted  by  judicious  misplacing  of 
the  knife-edges.  For  if  the  pan  knife-edges  be  placed  so  as  to  lie 
in  a  plane  slightly  above  the  plane  of  the  middle  knife-edge  when 
the  pans  carry  no  load  the  flexure  of  the  beam  when  under  the 
maximum  load  may  serve  to  bring  them  into  the  plane  of  the 
middle  knife-edge,  and  the  sensitiveness  of  the  instrument  under 
heavy  loading  will  be  thus  improved. 

In  a  perfect  balance  the  distance  between  the  beam  knife-edge 
and  the  two-pan  knife-edges  should  be  equal,  in  order  that  the 
leverage  of  equal  masses  placed  upon  the  pans  may  exert  equal 
effects  upon  the  oscillation  of  the  beam,  but  this  is  a  condition 
very  difficult  to  fulfill  absolutely  under  varying  conditions  of 
temperature.  It  is  possible,  however,  as  will  be  shown,  to 
determine  and  eliminate  the  effect  of  inequality  in  the  arms  of 
the  balance. 

Determination  of  the  Zero-Point .  —  The  balance-case  having 
been  leveled  by  means  of  supporting  screws,  and  the  pans 
lightly  dusted  with  a  fine  brush,  the  first  step  in  the  process  of 
weighing  is  the  adjustment  of  the  balance  to  a  suitable  equi- 
librium. With  the  eye  directly  in  front  of  the  middle  of  the 
scale,  the  pans  and  beam  are  released  by  a  slow  even  motion 
of  the  releasing  mechanism.  Should  the  beam  fail  to  oscillate 
when  set  free  it  may  be  put  in  motion  by  touching  it  gently 
with  the  rider  which  hangs  above  it  on  the  sliding  rider-rod.  The 
pan-supports  and  the  beam-support  of  some  balances  are 
moved  by  the  same  milled-head,  the  pans  being  released 
slightly  in  advance  of  the  beam;  sometimes  the  spring-support 


14 


QUANTITATIVE  CHEMICAL  ANALYSIS 


of  the  pans  and  the  beam-support  are  controlled  by  independent 
mechanisms.  When  the  pans  and  the  beam-supports  are  con- 
trolled separately  the  beam  begins  to  oscillate  more  readily  if  set 
free  after  the  lowering  of  the  pan-supports,  but  extreme  care 
should  always  be  taken  to  avoid  jarring  when  the  beam  knife- 
edge  comes  in  contact  with  the  bearing. 

The  exact  limits  to  which  the  pointer  moves  for  an  amplitude 
of  four  or  five  divisions  on  either  side  of  the  center  of  the  scale 
are  observed  for  an  uneven  number  of  swings  from  the  center, 
and  the  mean  of  each  set  is  noted;  then,  counting  the  excursions 
to  the  left  —  and  those  to  the  right  +  ,  half  the  algebraic  sum 
of  the  means  is  taken  as  the  center  of  oscillation  or  the  zero-point 
of  the  balance.  Suppose,  for  example,  successive  readings  to  be 
noted  and  recorded  as  follows : 


Divisions  to  left. 

Divisions  to  right. 

-  5 
-  4.6 
-  4.3 

+  5.2 

+  4.8 

sum   -13.9 
Mean-  4.63 

+10 
+  5 

Then  -  ':—= —  -  =  +0.18  and  the  zero-point,  or  point  of  equi- 
librium of  the  balance  without  load,  will  fall  0.18  of  a  division 
on  the  right  of  the  center.  When  the  balance  is  in  perfect 
adjustment  its  zero-point  should  coincide  with  the  middle-point 
of  the  scale,  and  if  the  displacement  of  the  zero-point  is  more 
than  a  division  of  the  scale  the  balance  should  be  readjusted. 
But  for  ordinary  weighing  it  is  generally  preferable  to  take 
slight  displacements  of  the  zero-point  into  account  in  the 
weighings  rather  than  to  disturb  the  beam  by  touch  or  by 
changes  of  temperature  due  to  the  insertion  of  the  hand  within 
the  balance  case. 

The  Operation  of  Weighing.  —  In  the  operation  of  weighing, 
the  material  to  be  weighed  is  usually  put  upon  the  left  pan  of 
the  balance  and  the  weights  required  to  bring  about  equilibrium 


WEIGHING  AND  MEASURING  15 

are  put  upon  the  right  pan,  down  to  the  limit  at  which  the  rider 
may  be  used  upon  the  right  arm  of  the  beam,  precaution  being 
taken  never  to  put  corrosive  substances  directly  upon  the 
balance-pan.  Crucibles,  dishes,  weighing  bottles,  pieces  of 
massive  metal,  etc.,  may  be  placed  upon  the  pan,  but  ordinary 
materials  for  analysis  are  usually  weighed  in  an  exactly  counter- 
poised watch-glass,  or  in  a  suitable  weighed  container  such  as 
a  crucible  or  thin  glass  weighing-tube,  the  weight  of  the  material 
being  shown  by  the  difference  between  the  weights  of  the 
container  with  and  without  the  material.  With  the  material 
placed  upon  the  left  pan  of  the  balance  the  counterpoising  is 
begun  by  the  systematic  trial  of  weights  upon  the  right  pan  in 
the  order  of  their  arrangement  in  the  box.  It  is  easy  to  find  the 
lightest  single  weight  which,  when  placed  upon  the  middle  of 
the  right  pan,  makes  the  pointer  move  to  the  left  upon  cau- 
tiously releasing  the  beam  and  pans.  With  balances  which  have 
separate  pan-supports,  about  a  gram  in  excess  of  the  weight 
necessary  to  produce  equilibrium  is  needed  to  move  the  pan, 
so  that  weights  down  to  the  limit  of  a  gram  may  be  found 
without  releasing  the  pan-supports.  After  finding  the  lightest 
single  weight  which  is  just  too  heavy,  this  weight  is  returned  to 
the  box  and  replaced  by  the  next  smaller  weight  in  the  box,  and 
this  second  weight  is  supplemented  by  weights  taken  in  order, 
with  necessary  replacements  and  additions,  until  the  limit  is 
reached  at  which  the  rider  may  be  used  upon  the  beam.  This 
limit  may  be  5,  6,  10,  or  12  mgrm.,  according  to  the  division 
of  the  beam.  Thereafter,  the  rider  is  used  to  indicate  milli- 
grams and  fractions  of  a  milligram.  Throughout  the  process, 
unused  weights  are  returned  to  their  places  in  the  box,  so  that 
the  first  reading  of  the  weight  used  may  be  made  from  the 
vacancies  in  the  box  and  the  reading  checked  by  noting  the 
weights  themselves  as  they  are  taken  subsequently  from  the  pan 
and  returned  to  their  places  in  the  box. 

The  Method  of  Equal  Swings.  —  A  simple  method  of  weigh- 
ing and  one  sufficiently  accurate  for  ordinary  analytical  proc- 
esses is  the  method  of  equal  swings.  If  the  balance  is  so 
adjusted  that  the  zero-point  is  displaced  by  half  the  amount 


16 


QUANTITATIVE  CHEMICAL  ANALYSIS 


lost  on  each  swing  from  left  to  right  and  back  to  left  the  pointer 
will  stop  at  the  end  of  each  swing  to  the  right  on  the  division 
which  corresponds  to  the  starting  point  at  the  left.  Then,  if 
the  weights  used  are  so  taken  that,  at  the  end  of  each  swing  to 
the  right,  the  pointer  stops  on  the  division  corresponding  to  the 
starting  point  on  the  left  the  amount  of  the  weights  as  read 
will  need  no  correction. 

Suppose,  for  example,  that  the  pointer  loses  0.3  of  a  division 
on  each  swing  of  moderate  amplitude  from  left  to  right  and  the 
reverse.  Then,  if  the  balance  is  adjusted  so  that  the  zero-point 
shall  fall  at  0.15  of  a  division  to  the  right  of  the  center,  the 
pointer  will  stop  in  successive  excursions  to  the  right  at  a  divi- 
sion which  corresponds  to  the  starting  point  on  the  left.  This 
is  shown  in  the  following  example. 


Distance  from  the 
zero-point  on  the  left. 

Zero-point. 

Distance  from  the 
zero-point  on  the  right. 

Loss  on  the  swing 
from  left  to  right  and 
the  reverse. 

-6.15 

0 

+6.0 

-5.85 

0 

+5.7 

0.30 

-5.55 

0 

+5.4 

0.30 

Starting-point  on  the 
left. 

Zero-point. 

Stopping-point  on  the 
right. 

-6.0 
-5.7 
-5.4 

+0.15 
+0.15 
+0.15 

+6.0 

+5.7 
+5.4 

When  the  weighing  is  made  by  the  method  of  equal  swings 
any  necessary  correction  for  a  known  imperfect  adjustment  of 
the  zero-point  may  be  made  mentally  by  adding  to,  or  sub- 
tracting from,  the  weights  used,  as  the  case  may  be,  the  weight 
which  would  be  necessary  to  adjust  the  zero-point. 

The  Method  of  Oscillations.  —  In  processes  demanding  the 
most  exact  weighing  the  equilibrium  is  found  just  as  in  the  de- 
termination of  the  zero-point  of  the  unloaded  balance,  by  a 
method  of  oscillations.  If  the  point  of  equilibrium  under  the 
load  and  the  counterpoising  weights  were  to  be  determined  in 


WEIGHING  AND  MEASURING  17 

the  same  way  that  the  zero-point  was  determined,  and  made  to 
coincide  exactly  with  the  zero-point,  the  weights  used  would 
indicate  at  once  the  weight  of  the  load.  This  process,  however, 
is  unnecessarily  tedious.  It  is  more  convenient  to  determine  a 
point  of  equilibrium  not  very  far  removed  from  the  zero-point  - 
and  then  to  make  a  correction  for  the  deviation;  and  this  may 
be  readily  done,  provided  the  sensitiveness  of  the  balance  (that  is, 
the  weight  which  corresponds  to  a  deviation  of  one  scale  divi- 
sion) is  known.*  This  correction  is  to  be  added  when  the  point 
of  equilibrium  under  the  load  lies  to  the  right  of  the  zero-point 
and  to  be  subtracted  when  the  point  of  equilibrium  lies  to  the 
left  of  the  zero-point.  For  example :  Suppose  the  weights  used, 
the  zero-point  of  the  unloaded  balance,  the  point  of  equilibrium 
of  the  loaded  balance,  and  the  value  of  one  scale  division  to  be 
as  given  in  the  tabular  statement  following. 

Weights  used,  grm.  Scale  divisions. 

20  Zero-point       +  0.2 

1  Point  of  equilibrium       —  0.4 

0-5  Deviation         .   0.6  to  the  left 

0.02  Value  of  1  scale  division  f  =  0.00043  grm. 

0.0057  (rider) 
21.6257 
Then, 

the  correction  for  deviation  =  0.6  X  0.00043  =  0.00026  grm. 
and  the  corrected  weight  of  load  =  21.6257  -  0.00026 

=  21.62544  grm. 

The  Sensitiveness  of  the  Balance.  —  To  find  the  sensitiveness 
of  the  balance  when  carrying  no  load  it  is  sufficient  to  note  the 
change  in  weight  indicated  by  the  rider  for  a  change  of  1  scale 
division  in  the  point  of  equilibrium;  and  the  sensitiveness  of  a 
loaded  balance  may  be  similarly  indicated  by  the  weight  which 
must  be  added  to,  or  subtracted  from,  the  counterpoising  weights 
in  order  to  change  the  point  of  equilibrium  by  1  scale  division. 
With  increasing  loads  the  friction  on  the  knife-edge  increases 
and  the  sensitiveness  of  the  balance  decreases.  So  it  is  impor- 
tant to  know  the  sensitiveness  for  various  loads.  With  any 
given  load,  or  no  load,  the  sensitiveness  may  be  found  by 
determining  the  variation  made  in  the  center  of  oscillations 
(the  point  of  equilibrium  or  the  zero-point)  by  moving  the  rider 
so  as  to  change  by  1  mgrm.  the  counterpoising  weights,  noting 

*  For  the  determination  of  the  sensitiveness  of  the  balance,  see  the  next 
paragraph. 
f  See  p.  18. 


18 


QUANTITATIVE  CHEMICAL  ANALYSIS 


the  effect  upon  the  point  of  equilibrium,  and  then  calculating 
the  weight  which  would  produce  a  variation  of  1  scale  division. 
Suppose,  for  example,  that  the  zero-point  has  been  found  to 
be  at  0.18  of  a  division  to  the  right  of  the  center,  and  that  the 
placing  of  the  rider  on  the  1-mgrm.  mark  will  cause  the  oscilla- 
tions shown  in  the  following  table. 


Divisions  to  left. 

Divisions  to  right. 

-  6 
-  5.4 
-  5.1 

+0.46 
+0.42 

Sum  -16.5 
Mean—  5.5 

+0.88 
+0.44 

Then  (-5.5  +  0.44)  -=-  2  =  -2.53  =  Point  of  equilibrium.  The 
addition  of  1  mgrm.  moves  the  point  of  equilibrium  *of  the 
balance  from  +0.18  to  —2.53  on  the  scale,  and  the  weight  of 
1  mgrm.  corresponds  to  2.7  divisions.  The  value  of  1  scale 
division  for  the  balance  without  load  will,  therefore,  be  found 
by  solving  the  proportion  2.7  :  1  =  1  :  x,  and  x  =  0.37  mgrm. 
The  sensitiveness  of  the  balance  for  varying  loads  may  be 
found  similarly.  Thus,  if  the  movement  of  the  rider  by  an 
amount  corresponding  to  a  change  of  1  mgrm.  in  weight  pro- 
duces variation  of  2.7  divisions  with  a  load  of  1  grm.,  2.7  divi- 
sion with  5  grm.,  2.6  divisions  with  10  grm.,  2.3  divisions  with 
20  grm.,  2.2  divisions  with  50  grm.,  these  figures  may  be  used  to 
find  the  sensitiveness  of  the  balance  under  loads  ranging  from 
1  grm.  to  50  grm.  It  is  convenient  to  tabulate  the  results  as 
follows  and  to  place  the  record  upon  a  card,  always  visible  in 
the  balance-case. 


Load. 

Value  of  scale 
division. 

grm. 

0 

1 

5 
10 
20 
50 

grm. 

0.00037 
0.00037 
0.00037 
0.00038 
0.00043 
0.00045 

In  ordinary  analytical  processes  the  determination  of  weights 
to  the  fourth  decimal  place  is  quite  as  exact  as  the  conditions 


WEIGHING  AND  MEASURING  19 

warrant.  If  the  fifth  decimal,  derived  by  calculation,  is  less 
than  5  it  should  be  dropped;  if  5,  or  more  than  5,  it  should  be 
replaced  by  1  added  to  the  fourth  decimal. 

Correction  for  Inequality  in  the  Arms  of  the  Balance.  —  The 
method  of  weighing,  according  to  which  the  material  to  be 
weighed  is  always  placed  on  the  same  pan,  gives  the  ratio  of 
weights  compared  in  successive  weighings  independently  of  the 
slight  inequality  in  arms  of  the  balance  which  may  be  due  to 
faulty  construction  or  to  existing  conditions  of  temperature, 
and  for  most  analytical  processes  this  method  is  sufficient. 
Sometimes,  however,  the  determination  of  weights  must  be 
absolute,  as,  for  example,  in  the  testing  of  an  untried  gram 
weight  against  a  standard  gram.  In  such  cases  the  effect  of 
inequality  in  the  lengths  of  the  arms  of  the  balance  must  be 
eliminated  by  the  method  of  weighing  by  substitution  or  by  the 
method  of  double  weighing. 

Weighing  by  Substitution.  —  If  the  material  to  be  weighed  is 
placed  upon  the  one  pan  of  the  balance  and  counterpoised 
exactly,  the  weights  which  will,  when  substituted  for  the 
material,  exactly  balance  the  counterpoise  will  be  the  exact 
measure  of  the  weight  of  the  material  independently  of  inequal- 
ity in  the  balance  arms.  In  this  process  it  is  more  convenient 
to  place  the  material  on  the  right  side  of  the  balance  so  that  the 
rider  may  be  used  in  the  final  adjustment  of  the  substituting 
weights.  The  material  on  the  right  pan  may  then  be  counter- 
poised by  any  suitable  material  placed  upon  the  left  pan,  but  the 
most  convenient  material  for  the  counterpoise  is  a  set  of  weights. 
The  counterpoised  material  is  removed  from  the  right  pan,  and 
replaced  by  exact  weights  until,  with  the  help  of  the  rider,  the 
balance  is  brought  to  an  equilibrium. 

Double  Weighing.  —  By  weighing  material  in  the  usual 
manner,  then  making  another  weighing  similarly  but  with  the 
material  and  weights  interchanged  upon  the  pans,  and  taking 
the  arithmetical  mean  of  the  sum  of  the  weights  obtained  in  the 
two  weighings,  the  effect  of  any  reasonable  inequality  in  the 
arms  of  the  balance  may  also  be  eliminated  sufficiently.  Theo- 
retically, it  is  the  geometrical  mean,  or  the  square  root  of  the 
product  of  the  two  weights,  which  gives  the  real  weight,  but  the 
arithmetical  mean  gives  a  sufficiently  near  approximation. 

Corrections  for  Buoyancy  of  the  Air.  —  In  weighings  made  by 
any  of  the  processes  described  both  weights  and  the  material 
weighed  are  affected  by  the  buoyancy  of  the  air,  and  such 
weighings  fail  to  give  a  perfect  comparison  between  weights  and 
material  by  just  the  difference  between  the  buoyant  effect 


20  QUANTITATIVE  CHEMICAL  ANALYSIS 

exerted  upon  the  weights  and  that  exerted  upon  the  material. 
When  the  amounts  of  material  weighed  are  as  small  as  they 
usually  are  in  analysis  the  difference  is  not  appreciable  for 
ordinary  purposes;  but  when  the  amounts  of  material  are 
considerable,  or  when  the  materials  are  weighed  in  a  large  con- 
tainer, account  must  be  taken  of  the  atmospheric  buoyancy  in 
order  that  the  weights  handled  may  be  known  with  a  reasonable 
degree  of  exactness.  The  effects  of  buoyancy  upon  materials 
and  weights,  measured  by  the  weights  of  air  displaced  by  each, 
may  be  determined  exactly  if  the  volume  of  each  is  known 
together  with  the  weight  of  a  cubic  centimeter  of  air  under  the 
prevailing  conditions  of  temperature,  pressure,  and  moisture. 
For  all  ordinary  work,  the  corrections  for  apparent  weights  and 
materials  may  be  determined  accurately  enough  from  the 
average  weight  of  a  cubic  centimeter  of  air  under  laboratory 
conditions  and  the  approximate  specific  gravities  of  materials 
and  weights.  The  average  weight  of  1  cm.3  of  air  may  be  taken 
as  0.0012  grm.  and  the  specific  gravity  of  the  weight,  or  the 
weights  which  fill  1  cm.3,  as  8.4  for  brass  and  21.5  for  platinum. 
The  corrections  to  be  subtracted  from  the  indicated  weight  are 
therefore  0.0012  -r-  8.4  =  0.00014  grm.  for  every  gram  of  brass, 
0.0012  -7-  21.5  =  0.00005  grm.  for  every  gram  of  platinum;  and 
the  correction  to  be  added  to  the  indicated  weight  will  be 
0.0012  grm.  -r-  sp.  gr.  of  material  for  every  gram.  Suppose,  for 
example,  that  a  mass  of  silver  placed  upon  the  left  pan  of  the 
balance  is  exactly  counterpoised  by  15.6875  grm.  upon  the  right 
pan.  From  this  indicated  weight,  which  is  too  large  by  the 
amount  of  the  buoyancy  of  the  air  exerted  upon  the  weights  and 
too  small  by  the  buoyancy  exerted  upon  the  silver,  must  be 
subtracted  for  the  brass  weights  15  X  0.00014  =  0.0024  grm., 
and  for  the  platinum  weights  0.68  X  0.00005  =  0.00003  grm., 
or  0.00243  grm.  in  all;  and  to  it  must  be  added  (15.6875  - 
0.00243)*  X  0.0012  -*-  10.15  (the  sp.  gr.  of  silver)  =  0.00179. 
The  indicated  weight  of  the  silver  is,  therefore,  0.00064  grm.  too 
large,  and  the  true  weight  is  15.68686  grm. 

It  is  obvious  that  in  a  very  exact  analysis  in  which  large 
amounts  of  material  are  involved  correction  for  the  buoyancy 
of  the  air  upon  weights  and  material  is  essential.  The  buoyant 
effect  upon  the  container  of  the  material  is  best  eliminated  by 
counterpoising  with  another  similar  container.  In  ordinary 
analysis,  in  which  the  amounts  of  material  used  rarely  exceed 

*  In  the  present  case,  the  subtraction  of  the  buoyancy  of  air  upon  the 
weights  at  this  point  is,  of  course,  insignificant. 


WEIGHING  AND  MEASURING  21 

a  gram,  and  generally  are  less,  the  correction  for  the  buoyancy 
of  air  is  superfluous. 

The  Analytical  Weights.  —  The  set  of  analytical  weights 
usually  consists  of  individual  weights  aggregating  100  grm~ 
beside  the  fraction  of  a  gram.  Generally  there  is  one  weight  of 
50  grm.,  one  of  20  grm.,  two  of  10  grm.,  one  of  5  grm.,  one  of 
2  grm.,  three  of  1  grm.,  all  of  brass;  and  smaller  weights  of 
platinum,  one  of  500  mgrm.,  one  of  200  mgrm.,  two  of  100  mgrm., 
one  of  50  mgrm.,  one  of  20  mgrm.,  two  of  10  mgrm.,  one  of  5 
mgrm.,  one  of  2  mgrm.,  and  three  of  1  mgrm.;  with  a  rider  of 
12  mgrm.,  10  mgrm.,  6  mgrm.,  or  5  mgrm.,  according  as  the 
beam  is  divided  between  the  middle  point  and  the  pan  knife-edge 
into  12, 10,  6,  or  5  large  divisions,  respectively.  In  practice,  the 
rider  takes  the  place  of  the  weights  smaller  than  its  own  weight. 
Weights  of  the  same  denomination;  the  two  10-grm.  weights, 
the  three  1-grm.  weights,  the  two  100-mgrm.  weights  and  the 
two  10-mgrm.  weights  should  be  distinguished  by  one,  two,  or 
three  dots  made  with  a  fine  round-pointed  punch,  and  should 
stand  in  the  box,  and  be  used  in  the  weighings,  in  the  order 
indicated  by  the  marks  of  identification. 

In  a  perfect  set  of  weights  the  1-grm.  weights  will  represent 
exactly  standard  grams,  the  multiple  weights  will  be  the  real 
multiples,  and  the  fractional  weights  the  real  fractions  of  the 
standard  gram  which  are  indicated  by  their  denominations. 
For  most  purposes  of  analysis  it  is  not  essential  that  the  gram 
weight  be  exactly  the  standard  gram  with  the  other  weights 
exact  multiples  and  fractions  of  the  standard  gram;  but  it  is 
essential  that  the  weights  throughout  the  series  represent  with 
a  high  degree  of  accuracy  the  exact  ratios  indicated  by  their 
designations.  Generally,  weights  of  first  quality,  when  new, 
are  accurate  enough  for  ordinary  work,  but  it  is  desirable,  and 
for  the  most  accurate  work  essential,  that  the  degree  of  error 
affecting  the  weights  be  known,  and  determined  from  time  to 
time. 

The  Testing  of  Weights.  —  In  testing  the  analytical  sets  it  is 
convenient  to  have  at  hand  a  small  set  of  testing  weights,  with 


22  QUANTITATIVE  CHEMICAL  ANALYSIS 

which  the  regular  analytical  weights  may  be  compared.  Such  a 
set  of  testing  weights  may  conveniently  consist  of  weights  of 
500  mgrm.,  200  mgrm.,  100  mgrm.,  50  mgrm.,  20  mgrm.,  10 
mgrm.,  and  5  mgrm.,  each  of  exactly  standard  weight  or  of 
known  error.  With  such  a  set  of  testing  weights  the  operation 
of  correcting  the  analytical  weights  by  the  method  of  substitu- 
tional  weighing  against  counterpoises  is  a  simple  matter. 

Let  us  suppose,  for  example,  that  the  standard  500-mgrm. 
weight  is  placed  upon  the  left  pan  of  the  balance  and  counter- 
poised by  weights  from  another  set  and  the  rider.  If  now  the 
standard  500-mgrm.  weight  is  replaced  by  the  500-mgrm.  weight 
of  the  set  to  be  corrected,  it  is  easy  to  determine,  by  moving 
the  rider  or  by  noting  the  variation  in  the  swing  of  the  pointer, 
just  how  the  tested  weight  compares  with  the  standard  500 
mgrm.  The  smaller  weights  may  be  compared  in  precisely  the 
same  manner  with  the  small  testing  weights.  A  similar  counter- 
poising of  the  first  gram  weight  of  the  set  and  substitution  by 
the  two  500-mgrm.  weights,  now  both  of  known  value,  will  fix 
the  value  of  that  weight;  and  similar  operations  with  the  second 
and  third  1-grm.  weights  will  determine  their  values  in  terms  of 
the  standard  gram;  and  for  the  weights  heavier  than  1-grm., 
the  corrections  may  be  determined  successively  by  substitu- 
tional  comparison  with  the  weights  already  corrected.  Thus, 
the  2-grm.  weight  may  be  compared  with  two  corrected  1-grm. 
weights;  the  5-grm.  weight  with  the  three  1-grm.  weights  and 
the  2-grm.  weight;  the  10-grm.  weight  with  the  5-grm.,  the  2- 
grm.,  and  the  three  1-grm.  weights;  and  so  on. 

In  case  such  a  set  of  small  standard  weights  is  not  at  hand, 
it  is  possible,  also,  to  correct  the  ratios  of  the  individual  weights 
of  a  set  by  systematic  comparisons  with  one  another  either  by 
the  method  of  double  weighing*  or  much  more  simply  by  sub- 
stitutional  weighing  against  counterpoises.!  Such  an  operation 
is  best  begun  by  counterpoising  exactly  (by  weights  and  rider) 
an  extra  weight  a  trifle  lighter  than  0.01  grm.  and  then  sub- 
stituting for  it  the  0.01 '-grm.  weight  of  the  set  to  be  tested, 

*  Ostwald:  Physico-Chemical  Measurements,  trans.  Walker,  p.  41. 
f  Richards:  Jour.  Am.  Chem.  Soc.,  22,  144  (1900). 


WEIGHING  AND  MEASURING 


23 


determining  the  relative  values  of  these  weights  by  means  of 
the  rider.  Next  may  be  found  similarly  the  relation  between 
the  0.01  "-grm.  weight  and  the  0.01  '-grm.  weight;  then  the  re- 
lation between  the  0.02-grm.  weights  and  the  0.01 '-grm.  and 
0.01  "-grm.  weights  together;  then  the  relation  between  the 
0.05-grm.  weight  and  the  sum  of  the  0.01  '-grm.,  0.01  "-grm., 
0.02/-grm.  and  the  extra  0.01-grm.  weights,  and  so  on.  In  this 
manner  a  statement  of  ratios  may  be  constructed,  like  that  given 
in  the  first  and  second  columns  of  the  following  table.  In  mak- 
ing use  of  a  small  weight  as  the  standard  of  comparison  the  errors 
of  the  large  weights  are  made  to  appear  very  large;  but  these 
relative  values  once  determined,  more  convenient  ratios  may 
be  found  by  making  a  large  weight  the  standard  and  redeter- 
mining  the  ratios.  For  example,  if  the  10-grm.  weight  is  now 
taken  as  the  standard,  the  ideal  aliquot  parts  of  it  represented  by 
the  smaller  denominations  will  be  shown  in  the  third  column, 
and  the  difference  between  these  ideal  aliquot  parts  and  the  cor- 
responding values  of  the  second  column  will  measure  the  errors 
of  the  individual  weights. 


Denomination  of 
weight. 

Ratios  with  the  extra 
0.01-grm.  weight  as 
a  standard. 

Weights  of  each  denom- 
ination calculated  from 
the  10-grm.  weight 
taken  as  a  standard. 

Correction. 

Grm. 

Grm. 

Grm. 

Grm. 

0.01 
0.01 

0.01006) 

0.01005  J 

0.01002 

j  +0.00004 

I  +0.00003 

0.02 

0.02005 

0.02004 

+0.00001 

0.05 

0.05009 

0.05009 

+0.00000 

0.1 
0.1 

0.10019  ) 
0.100201 

0.10018 

j  +0.00001 
1  +0.00002 

0.2 

0.20035 

0.20035 

±0.00000 

0.05 

0.50088 

0.50088 

±0.00000 

1 

1.00183) 

(  +0.00006 

1 

1.00181  > 

1.00177 

<  +0.00004 

1 

1.00177) 

(  +0.00000 

2 

2.00383 

2.00354 

+0.00029 

5 

5.00884 

5.00884 

±0.00000 

10 

10.01768 

10.01768 

Standard 

In  making  use  of  such  a  table  of  corrections,  +  corrections  are, 
of  course,  to  be  added  to  the  respective  weights  and  —  correc- 
tions subtracted. 

Generally,  analytical  processes  deal  with  amounts  of  materials 
which  do  not  weigh  more  than  a  gram  or  two,  the  higher  weights 
being  used  simply  as  counterpoises  of  crucibles  or  other  con- 


24  QUANTITATIVE  CHEMICAL  ANALYSIS 

tainers.  For  such  processes,  therefore,  it  is  not  necessary  to 
use,  or  even  to  know,  the  true  values  of  the  higher  weights, 
provided  identical  counterpoises  are  used  throughout  a  process. 

The  Measurement  of  Volume 

Volumetric  Apparatus.  —  In  order  that  measurements  of 
volumes  of  liquids  and  gases,  for  which  volumetric  measure- 
ments are  feasible,  may  be  exact,  it  is  obvious  that  accurately 
graduated  and  strictly  comparable  measuring  vessels  are 
requisite.  At  the  present  time  measuring  instruments  are  pro- 
curable which  are  certified  by  responsible  governmental  au- 
thority.* These  instruments  are  graduated  with  reference  to 
the  liter  defined  as  the  volume  occupied  by  one  kilogram  of  pure 
water  at  4  degrees,  the  temperature  of  its  greatest  density;  and 
this  is  the  system  of  graduation  to  which  it  is 
advisable  that  all  instruments  employed  in  the 
measurements  of  volumetric  analysis  should  con- 
form. The  measuring  vessels  commonly  used  in 
volumetric  analysis  are  flasks,  Fig.  2,  graduated 
for  contents  or  for  delivery;  single-volume  pipettes, 
Fig.  3a  (for  the  delivery  of  one  definite  volume) ; 
and  divided  pipettes,  Fig.  3b,  without  stop-cocks; 
burettes,  Fig.  4a  and  Fig.  46,  provided  with  stop- 
cocks, or  pinch-cocks,  or  with  the  Bunsen  tap  (con- 
sisting  of  a  bead  within  a  rubber  tube,  which  by 
gentle  manipulation  between  the  thumb  and  fore- 
finger allows  the  passage  of  liquid  in  regulated  amount),  for 
the  delivery  of  various  volumes  throughout  the  range  of 
graduation;  and  gas-burettes,  Fig.  5,  fitted  with  stop-cocks  and 
leveling  tubes  for  the  adjustment  of  pressure. 

It  is,  of  course,  necessary  in  all  exact  processes  that  readings 
should  be  made  with  the  eye  on  a  level  with  the  reading  mark 
in  order  to  avoid  parallax,  and  the  graduation  of  the  tube  by 
lines  entirely  encircling  it  is  the  best  aid  to  a  correct  placing  of 
the  eye  in  reading.  In  determining  the  volumes  of  liquids  the 

*  Die  Physikalische-Technische  Reichsanstalt  of  Germany  and  the 
National  Bureau  of  Standards  of  the  United  States. 


WEIGHING  AND  MEASURING 


25 


reading  is  made  from  the  upper  boundary  when  they  are  opaque 
and  preferably  from  the  bottom  of  the  meniscus  when  they  are 
transparent.  In  the  former  case,  a  white  card  held  behind  the 
graduation  helps  to  define  the  reading  boundary;  in  the  latter, 
a  white  card  striped  horizontally  with  a  broad  band  of  black 


a  b 

FIG.  3. 


FIG.  4. 


FIG.  5. 


and  held  so  that  the  darkened  meniscus  may  be  defined  clearly 
is  an  efficient  aid. 

Variations  of  Volume  with  Changing  Temperature.  —  Gases 
expand  and  contract  regularly  with  changes  in  temperature  and 
volumes  of  gases  measured  at  known  temperatures  and  pressures 
may  be  reduced  to  normal  or  definite  conditions  by  application 
of  the  laws  of  Boyle  and  Charles. 

Liquids  and  solutions  expand  or  contract  with  changes  in 
temperature  by  amounts  which  vary  with  different  ranges  of 
temperature  as  well  as  with  the  nature  and  concentrations  of 


26  QUANTITATIVE  CHEMICAL  ANALYSIS 

the  substances.  For  example,  a  liter  of  pure  water  expands 
2.05  cm.3  in  the  range  of  10  degrees  from  15  degrees  to  25  degrees, 
and  2.55  cm.3  in  the  range  from  20  degrees  to  30  degrees.  In 
the  range  from  15  degrees  to  25  degrees,  the  expansion  of  a  liter 
of  a  solution  containing  3.16  grm.  of  potassium  permanganate* 
amounts  to  2.13  cm.3,  and  that  of  a  liter  of  a  solution  containing 
17  grm.  of  silver  nitrate  f  amounts  to  2.16  cm.3;  that  of  a  liter 
of  hydrochloric  acid  containing  36.5  grm.  of  that  reagent  | 
amounts  to  2.42  cm.3,  and  that  of  a  liter  of  a  solution  con- 
taining 46  grm.  of  sodium  hydroxide  {  amounts  to  3.15  cm.3. 

The  measuring  vessels  are  also  subject  to  variations  of  volume 
with  changing  temperatures.  In  the  average,  glass  vessels 
expand  or  contract  by  about  0.000026  of  their  volume  for  every 
degree  of  rise  or  fall  in  temperature,  respectively.  The  expan- 
sion or  contraction  of  a  liter  flask  for  a  range  of  10  degrees  may 
be  taken  as  1000  X  10  X  0.000026  =  0.26  cm.3. 

It  is  plain,  therefore,  that  graduated  apparatus  should  be 
used  for  the  determination  of  volumes  of  a  solution  at  a  tem- 
perature not  very  different  from  the  temperature  of  standard- 
ization, usually  15  degrees,  unless  correction  is  made  for'  the 
variations  of  the  volumes  with  changing  temperature. 

Correction  of  the  Indicated  Volume.  —  For  change  in  the 
volume  of  the  apparatus,  the  average  expansion  of  glass  for  a 
rise  of  1  degree  (the  coefficient  of  expansion)  and  the  range  of 
temperature  afford  the  data  necessary  to  make  due  correction; 
and  a  correction  of  the  indicated  volume  for  change  in  the 
solution  itself  may,  if  the  coefficient  of  expansion  of  the  solution 
be  known,  be  similarly  applied.  The  coefficients  of  expansion 
of  very  dilute  solutions  generally  do  not  vary  much  from  that 
of  pure  water.  For  such  solutions  containing  only  water  and 
the  solute  at  very  low  concentration,  the  corrected  volume  V 
may  be  found  by  adding  to  the  observed  volume  V  the  product 
obtained  by  multiplying  the  observed  volume  by  a  factor  a 
depending  upon  the  temperature  of  observation,  this  factor 
bearing  the  +  sign  for  observations  made  below  15  degrees 
(the  ordinary  temperature  of  graduation)  and  the  —  sign  for 
observations  made  above  15  degrees.  That  is,  V  =  V  -f  Va. 

*  0.02  grm.  mol.  f  0-1  Srm.  mol.  |  1  grm.  mol. 


WEIGHING  AND  MEASURING 


27 


The  approximate  values  of  this  factor  for  solutions  containing 
0.1  grm.  mol.  of  a  single  reagent  in  1  liter  are  shown  in  the  fol- 
lowing table. 


Temperature    

10° 

11° 

12° 

13° 

14° 

15° 

Factor  =  -f-0  001  X  

0.46 

0.40 

0.33 

0.22 

0.12 

0 

Temperature 

16° 

17° 

18° 

19° 

20° 

Factor  —  —  0  001  X 

0.13 

0  37 

0.42 

0.59 

0.76 

Temperature 

21° 

22° 

23° 

24° 

25° 

Factor  —  —0  001  X  . 

0.95 

1.14 

1.35 

1.56 

1.79 

Temperature 

26° 

27° 

28° 

29° 

30° 

Factor  —  —  0  001  X 

2.02 

2.27 

2.52 

'2.75 

3.06 

For  example,  according  to  a  computation  made  by  the  aid  of 
this  table,  a  volume  of  50  cm.3  of  a  0.1  grm.  mol.  solution  of  a 
reagent  (and  no  other  dissolved  material)  if  measured  at  25 
degrees,  would  measure 

50  -  50  X  0.001  X  1.79  =  49.91  cm.3  at  15  degrees; 
and  a  volume  of  50  cm.3  at  10  degrees  would  measure 

50  -f  50  X  0.001  X  0.46  =  50.02  cm.3  at  15  degrees. 

Frequently,  however,  solutions  are  employed  of  much  greater 
concentration  in  respect  to  the  active  reagent  or  to  all  the  dis- 
solved substances;  and  solutions  are  often  <ss^, 
standardized  at  a  temperature  different  from 
the  temperature  of  graduation.  In  general,  J\ 
therefore,  corrections  of  volume  for  varying 
temperatures  of  solutions  may  be  based  pref- 
erably upon  an  experimental  determination 
of  the  ratio  of  the  concentration  of  the  solu- 
tion at  the  temperature  of  standardization  to 
the  concentration  at  any  temperature  of  use. 
The  Sprengel  pyknometer  (Fig.  6)  as  modified 
by  Ostwald  *  serves  the  purpose. 

The  pyknometer,  of  about  10  cm.3  capacity, 
washed  with  water,  alcohol,  and  ether,  applied  successively,  and 
dried  by  drawing  air  through  it,  is  first  weighed  empty.     Next, 


FIG.  6. 


*  Ostwald:  Physico-Chemical  Measurements,  trans.  Walker,  p.  92. 


28  QUANTITATIVE  CHEMICAL  ANALYSIS 

the  pyknometer  is  filled  with  the  solution  at  the  temperature  of 
standardization;  the  volume  of  the  solution  is  adjusted  to  the 
mark  by  touching  the  narrowed  end  of  the  tube  with  filter-paper 
to  withdraw  liquid,  or  with  a  drop  of  solution  on  a  rod  to 
introduce  liquid;  and  the  pyknometer  and  solution  are  weighed. 
The  difference  between  the  weight  of  the  empty  instrument  and 
the  full  instrument  is  the  weight  of  a  definite  volume  of  the 
solution  at  the  temperature  of  standardization.  Similarly,  the 
difference  between  the  weight  of  the  empty  pyknometer  and 
the  weight  of  the  pyknometer  filled  to  the  mark  with  the  solution 
at  the  temperature  of  use  gives  the  weight  of  a  pyknometerful 
of  the  solution  for  the  two  temperatures. 

Suppose,  for  example,  that  a  solution  of  silver  nitrate  is  made 
to  contain  17  grm.  of  that  reagent  in  the  liter,  or  1.7  grm.  in 
100  cm.3,  at  20  degrees,  and  that  the  pyknometer  contains 
10.1780  grm.  at  20  degrees  and  10.1680  grm.  at  25  degrees,  then 
the  amount  of  silver  nitrate  contained  in  the  volume  indicated 
;as  100  cm.3  by  a  glass  measuring  instrument  at  25  degrees  will 
be  given  by  the  proportion 

10.1780  :  10.1680  =  1.7  :  x,        x=  1.6983  grm. 

Calibration  of  Instruments.  —  It  is  often  desirable  to  calibrate 
graduated  instruments,  or  to  graduate  less  expensive  apparatus, 
and  this  may  be  accomplished  by  means  of  the  relation  which 
exists  between  the  weight  of  water  and  the  volume  which  it 
occupies  at  given  temperatures. 

In  weighing  the  water  it  is  necessary  to  take  into  considera- 
tion the  buoyancy  of  the  air  both  upon  the  brass  weights  and 
upon  the  corresponding  mass  of  water;  but  weights  of  less  than 
1  grm.  may  be  safely  left  out  of  account.  The  average  buoy- 
ancy exerted  upon  1  grm.  of  brass  amounts  to  0.00014  grm.  and 
that  exerted  upon  1  grm.  of  water  to  0.0012  grm.  To  obtain 
the  true  weight  of  the  water  weighed  in  air,  the  sum  of  the 
weights  actually  used  must  be  increased  by  0.00106  grm.,  the 
difference  between  0.0012  grm.  and  0.00014  grm.,  for  every 
gram.  On  the  other  hand,  to  find  the  weight  which  must  be 
used  in  air  to  counterpoise  an  amount  of  water  which  weighs  a 
given  number  of  grams  in  a  vacuum,  the  given  weight  must  be 
diminished  by  0.00106  grm.  for  every  gram. 

The  following  table  gives  the  actual  weights  (in  a  vacuum) 
of  1  cm.3  of  water  at  ordinary  laboratory  temperatures  and  the 
weights  which  must  be  used  to  counterpoise  in  air  the  actual 
weights. 


WEIGHING  AND  MEASURING 


29 


Temperature, 
Deg.  C. 

Weight  of  1  cm."  of 
pure  water  in  a 
vacuum. 

Counterpoising  weights 
in  air  for  1  cm.3  of 
water. 

4 

Grm. 

1.000000 

Grm. 

0.998940 

15 

0.999126 

0.998066 

16 

0.998970 

0.997910 

17 

0.998801 

0.997741 

18 

0.998622 

0.997562 

19 

0.998432 

0.997372 

20 

0.998230 

0.997170 

21 

0.998019 

0.996959 

22 

0.997797 

0.996737 

23 

0.997565 

0.996505 

24 

0.997323 

0.996263 

25 

0.997071 

0.996011 

26 

0.996810 

0.995750 

27 

0.996539 

0.995479 

28 

0.996259 

0.995199 

29 

0.995971 

0.994911 

30 

0.995673 

0.994613 

These  relations  of  the  weights  of  pure  water  to  the  volumes  make 
it  easy  to  determine  or  adjust  the  accuracy  of  measuring  appara- 
tus at  any  convenient  temperature. 

Flasks.  —  The  calibration  or  testing  of  a  flask  graduated  for 
contents  may  be  accomplished  by  counterpoising  the  dry  flask, 
filling  it  with  water  until  the  bottom  of  the  meniscus  of  liquid 
touches  the  mark  upon  the  neck,  finding  the  weight  of  the  flask 
with  the  water,  and,  by  difference,  the  weight  of  the  water,  and 
then  comparing  the  apparent  weight  of  the  water  with  the 
actual  weight  in  air  of  the  indicated  volume  of  water  for  the 
temperature  at  the  time.  For  example,  if  the  weight  of  water 
required  to  fill  to  the  mark  a  liter  flask  proves  to  be  997.37  grm. 
at  20  degrees,  the  error  of  the  graduation  will,  according  to  the 
table,  be  the  difference  between  this  weight  and  997.170  grm., 
the  weight  in  air  of  1000  cm.3  of  water  at  20  degrees,  or  +0.20 
cm.3. 

In  calibrating  a  flask  for  delivery  it  is  first  weighed  full  of 
water  to  the  mark  and  then  emptied,  drained  in  an  inclined 
position  for  a  definite  period  (one  minute)  and  weighed  again. 
The  difference  between  the  weights  noted  measures  the  contents 
of  the  flask  for  delivery.  In  using  the  flask  for  measurements 
it  is  allowed  to  drain  for  the  same  definite  period. 

Ungraduated  flasks  may  be  graduated  by  counterpoising  dry 
or  after  draining,  according  as  they  are  to  be  used  for  contents 
or  for  delivery,  and  then  filling  them  with  exactly  the  amount 


30  QUANTITATIVE  CHEMICAL  ANALYSIS 

of  water  which  has  the  weight  of  the  desired  volume  of  water 
at  the  given  temperature,  and  marking. 

Pipettes.  —  For  the  delivery  of  liquids  in  exact  work  cali- 
brated pipettes  are  preferable  to  flasks.  The  pipette  is  filled 
by  suction  until  the  bottom  of  the  meniscus  of  liquid  is  a  little 
above  the  mark  upon  the  stem  and  closed  at  the  top  with  the 
finger.  By  moving  the  finger  across  the  tube  held  in  a  vertical 
position  and  with  the  lower  end  touching  the  side  wall  of  a  glass 
beaker,  the  meniscus  is  adjusted  exactly  to  the  mark.  In  using 
a  single-volume  pipette  the  liquid  is  allowed  to  run  out  entirely 
into  a  counterpoised  flask,  and  with  the  point  in  contact  with 
the  neck  of  the  flask  the  pipette  is  drained  for  fifteen  seconds. 
From  divided  pipettes  the  liquid  is  run  into  the  flask  to  the  de- 
sired point  and  drained  for  fifteen  seconds,  and  the  meniscus 
is  again  adjusted.  The  increase  in  the  weight  of  the  flask 
measures  the  volume  of  the  water  delivered. 

In  graduating  unmarked  pipettes,  the  upper  point  of  marking 
is  first  determined  approximately  by  filling  the  pipette,*  from 
the  orifice  or  from  a  fixed  mark,  as  the  case  may  be,  with  a  vol- 
ume of  water  slightly  larger  than  that  for  which  the  graduation 
is  to  be  made,  and  the  level  of  the  meniscus  is  marked  upon  a 
paper  scale  pasted  upon  the  stem.  Then  two  weighings,  one  of 
the  volume  delivered  from  a  noted  point  above  the  mark,  and 
one  from  a  noted  point  below  the  mark,  to  the  orifice  or  to  the 
fixed  mark,  will  give  the  data  for  determining  the  exact  point  of 
graduation.  For  example,  if  one  weighing  of  water,  at  18  de- 
grees, from  a  pipette  designed  to  deliver  10  cm.3  proves  to  be 
10.032  grm.  and  the  other  from  a  point  10  mm.  lower  proves 
to  be  9.892  grm.,  then,  10.032  -  9.892  :  10.032  -  9.976  (the 
weight  of  10  cm.3  of  water  in  air  at  18  degrees)  =  10  :  x,  and 
x  =  4;  and  the  graduation  mark  should  be  made  at  the  point 
4  mm.  below  the  uppermost  mark.  If  the  weighings  were  made 
at  20  degrees,  the  mark  of  graduation  should  be  placed  4.3  mm. 
below  the  uppermost  point,  since  10.032  -  9.892  :  10.032  - 
9.972  =  10  :  x,  and  x  =  4.3  mm. 

Burettes  and  Divided  Pipettes.  —  Burettes  for  the  measure- 
ment of  liquids  or  gases,  and  the  divided  pipette,  may  be  tested 
by  weighing,  to  0.01  grm.,  carefully  measured  portions  of  water 
drained  successively  into  a  counterpoised  flask,  and  comparing 
the  weights  of  the  indicated  volumes  with  the  weights  which 
such  volumes  should  have  at  the  temperatures  of  the  experiment. 
The  corrections  of  readings  once  determined  may  be  made  easily 
available  for  practical  application  by  plotting  the  readings  as 
abscissae  and  the  corrections  as  ordinates,  so  that  the  correction 
for  any  reading  may  be  seen  at  a  glance. 


WEIGHING  AND  MEASURING 


31 


A  more  convenient  method  of  testing  burettes  and  divided 
pipettes  involves  the  use  of  a  small  pipette,  conveniently  of 
5  crn.3  capacity  and  calibrated  in  the  manner 
previously  described.*  The  pipette  is  at- 
tached to  the  instrument  to  be  tested,  in  the 
manner  shown  in  the  accompanying  figure, 
and  the  water-levels  are  adjusted  to  the  zero- 
points.  Then  with  the  flow  cut  off  from  the 
instrument  to  be  tested,  the  water  is  drawn 
from  the  calibration  pipette  to  the  lower 
mark,  the  cock  is  closed,  and  the  liquid 
allowed  to. run  from  the  tested  instrument 
into  the  calibration-pipette  until  the  latter 
is  again  filled  to  the  upper  mark.  The 
difference  between  the  reading  of  the  tested 
instrument  at  the  start  and  the  reading  at 
the  end  of  this  operation  will  show  the 
volume  which  for  the  temperature  of  the 
experiment  equals  the  capacity  of  the  cali- 
bration pipette.  In  this  manner  the  entire 
instrument  may  be  calibrated.  When  much 
calibration  of  apparatus  is  to  be  done  it  is 
most  convenient  to  make  use  of  comparison 
pipettes  with  stems  graduated  and  calibrated  to  deliver  exact 
volumes  for  any  convenient  temperature.!  The  limits  within 
which  errors  of  graduation  may  be  regarded  as  reasonable  may 
be  stated  as  follows  :| 


FIG.  7. 


Cm.3 

Cm." 

10 

0.01 
0.03 

0.02 
0.04 
0.02 

Cm." 

50 

0.05 
0.10 

0.05 
0.10 
0.05 

Cm.3 

Cm." 

Cm." 

Cm.' 

Cm.' 

Capacity  
Limit  of  error,  flasks: 
for  contents 

1  to  2 

100 

0.08 
0.15 

0.10 
0.20 
0.10 

250 

0.10 
0.20 

500 

0.15 
0.30 

1000 

0.30 
0.50 

2000 

0.50 
1.00 

for  delivery 

Limit   of  error,    undivided  pi- 
pettes 

0.01 
0.02 

Limit  of  error,  divided  pipettes. 
Limit  of  error,  burettes  

*  Ostwald:   Physico-Chemical  Measurements,  trans.  Walker,  p.  87. 

f  Morse  and  Blalock:  Amer.  Chem.  Jour.,  16,  479. 

t  Further  discussion  of  specifications,  marks,  graduation,  precautions 
in  use,  and  allowable  error  in  volumetric  apparatus  may  be  found  in  the 
report  of  the  Committee  for  Cooperation  with  the  National  Bureau  of 
Standards.  Jour.  Amer.  Chem.  Soc.,  26;  Proc.,  p.  (17),  1904. 


CHAPTER  III 
PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 

In  processes  of  Gravimetric  Analysis  both  the  substances  ana- 
lyzed and  the  products  of  analysis  are  determined  by  weighing. 
When  exposed  to  ordinary  air,  some  substances  prove  to  be 
hygroscopic  (that  is,  they  take  up  moisture  from  the  air) ;  others 
give  up  water  to  ordinary  air;  while  others  assume  fairly  definite 
conditions  when  exposed  in  an  ordinarily  dry  atmosphere*  The 
water  content  of  every  substance  exposed  to  air  and  capable  of 
taking  up  or  of  losing  water  depends  upon  an  equilibrium 
between  the  vapor  pressure  of  the  water  in  the  substance  and 
that  of  the  water  in  the  air,  and  the  condition  of  equilibrium 
varies  with  temperature  and  pressure.  It  is  necessary,  there- 
fore, to  make  sure  that  the  substances  determined,  and  generally 
the  material  subjected  to  analysis,  ordinarily  weighed  in  contact 
with  air,  shall  be  definite  in  respect  to  content  of  moisture;  and 
the  means  of  attaining  this  end,  by  regulation  of  the  condition 
of  temperature  and  pressure,  must  be  chosen  with  a  knowledge 
of  the  special  characteristics  of  the  substances. 

Preparation  for  Weighing 

Air-Drying.  —  Some  substances  may  be  brought  to  a  fairly 
definite  condition  by  simple  air-drying  —  that  is,  by  exposing 
them  to  ordinary  air  for  a  suitable  period  at  the  ordinary  atmos- 
pheric temperature.  Of  such  substances,  crystallized  barium 
chloride,  BaCl2  •  2  H20,  and  crystallized  copper  sulphate, 
CuS04  •  5  H20,  which  lose  hygroscopic  water  while  retaining 
their  water  of  crystallization  in  ordinarily  dry  air,  may  serve 
as  examples. 

32 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


33 


Drying  in  Desiccators.  —  Some  substances  are  best  prepared 
for  weighing  by  exposing  them  in  closed  chambers,  at  the 
ordinary  temperature,  to  air  which  is  artificially  dried  by 
contact  with  some  strongly  hygroscopic  substance,  like  calcium 
chloride,  concentrated  sulphuric  acid,  or  phosphorus  pentoxide; 
so  that  the  water  of  the  substance  may  come  to  an  equilibrium 
with  water  vapor  at  inappreciable  concentration  in  the  drying 
atmosphere.  Sometimes  to  promote  the  drying  effect  the  air 
pressure  and  vapor  pressure  are  reduced  in  the  chamber  by 
means  of  the  vacuum  pump.  Many  forms  of  these  desiccators 
are  in  common  use,  but  the  essential  characteristics  of  every 


desiccator  are  the  closed  chamber  and  a  dehydrator  suited  to 
the  occasion,  and  in  certain  cases,  means  for  exhausting  the 
chamber  (Fig.  8). 

Dehydration  by  Heat.  —  To  bring  some  substances  to  a 
definite  condition  of  hydration  or  dehydration,  it  is  necessary 
to  change  the  condition  of  equilibrium  by  heating  them  in  the 
hot  air  chamber  (the  air-bath),  or  by  " igniting"  them  —  that  is, 
heating  them  in  the  flame  of  the  Bunsen  burner  or  the  blast-lamp 
or  other  source  of  a  high  degree  of  heat,  as  the  case  may  require ; 
and  when  a  substance  has  been  properly  dehydrated  by  heat  it 
is  usual  to  place  it  in  a  desiccator  in  order  that  it  may  be  pro- 
tected against  the  possibility  of  taking  up  water  by  exposure  to 


34  QUANTITATIVE  CHEMICAL  ANALYSIS 

the  ordinary  atmosphere  while  cooling  completely  to  the  atmos- 
pheric temperature,  at  which  it  must  be  weighed.* 

PROCESSES  DEPENDING  UPON  THE  FORMATION 
OF  VOLATILE  PRODUCTS 

Like  the  processes  of  dehydration  used  in  the  preparation 
of  substances  for  weighing,  some  procedures  of  gravimetric 
analysis  rest  upon  the  fixing  of  conditions  of  temperature  (and 
pressure)  so  that  the  substances  to  be  analyzed  shall  yield  a  volatile 
constituent  and  leave  a  definite  and  weighable  residue,  or  upon  the 
expulsion  of  the  volatile  constituent  by  a  reaction  resulting  in  the 
formation  of  a  definite  and  weighable  residue.  Other  gravimetric 
procedures  involve  the  collection  and  weighing  of  the  volatile  con- 
stituent, after  its  expulsion  from  the  substance  analyzed. 

The  Determination  of  Water  of  Composition 

The  water  content  of  certain  hydroxides  and  many  hydrous 
salts  may  be  expelled  by  simply  submitting  these  substances  to 
suitable  conditions  of  temperature  with  the  formation  of  a 
definite  and  weighable  residue.  The  difference  between  the 
weight  of  the  substance  taken  and  that  of  the  residue  will  then 
be  the  weight  of  the  water  evolved. 

From  certain  refractory  salts,  minerals,  and  rocks  from  which 
the  water  is  expelled  with  difficulty  by  heat  alone,  the  evolution 
of  the  water  may  be  facilitated  by  heating  with  a  suitable  flux 
of  unchanging  weight;  and  from  certain  other  substances  which 
when  heated  by  themselves  evolve  beside  the  water  content 
other  volatile  matter  the  water  may  be  evolved  alone  if  a  suitable 
flux  be  used  as  a  retainer  of  the  volatile  constituents  other  than 
water.  In  either  case  the  loss  of  weight  on  the  part  of  the 
mixture  of  substance  and  flux  will  measure  the  water  evolved. 

When  the  residue  left  after  heating,  either  with  or  without 
the  aid  of  a  flux,  is  not  sufficiently  definite  the  determination 
must  be  made  by  the  direct  weighing  of  the  evolved  water 
usually  collected  in  a  suitable  absorbent. 

*  For  the  description  of  a  special  device  by  which  substances  may  be 
dried  and  securely  bottled  without  contact  with  the  ordinary  air,  see 
Richards:  Methods  Used  in  Precise  Chemical  Investigation,  Publication 
No.  125  of  the  Carnegie  Institution  of  Washington,  p.  97. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


35 


Determination  of  Water  by  Loss  on  Heating.  —  Hydrous 
barium  chloride,  BaCl2  •  2  H2O,  and  hydrous  copper  sulphate, 
CuSOd  •  5  H2O,  both  crystallized  from  water  and  air-dried  under 
ordinary  atmospheric  conditions,  lose  their  water  of  crystalliza- 
tion when  suitably  heated  for  a  sufficient  length  of  time  under 
conditions  which  permit  the  escape  of  water  from  the  system. 
From  the  hydrous  barium  chloride  all  the  water  may  be  easily 
expelled  at  a  temperature  of  100  degrees.  To  remove  the  water 
entirely  from  the  hydrous  copper  sulphate  prolonged  heating 
at  a  much  higher  temperature  (200°  or  more)  is  required. 

Experimental  Process:  Determination  of  Water  in  Barium  Chloride.  — 
Clean  a  crucible  (porcelain  or  platinum),  ignite  it  in  the  Bunsen  flame, 
cool  it  until  it  radiates  heat  gently,  place  it  in  the  desiccator,  and,  after  it 
has  been  thoroughly  cooled  to  the  temperature  of  the  air  (at  least  ten  to 
fifteen  minutes)  place  it  upon  the  left-hand  pan  of  the  balance,  and  weigh 
carefully.  Put  into  the  crucible,  with  the  aid  of  a  spatula,  about  1  grm. 
of  barium  chloride,  BaCl2  •  2  H2O,  previously  prepared  by  powdering  in  a 
mortar  the  pure  air-dried  crystals.  Weigh  again  carefully.  The  difference 
between  the  two  weights  will  be  the  weight  of  the  substance  taken.  Now, 
place  the  crucible  in  an  air-bath  heated  at  115°  to  120°,  and  let  it  stay 
there  for  half  an  hour.*  Then  remove  the  crucible  to  the  desiccator,  let  it 
cool  for  fifteen  minutes,  and  weigh.  Repeat  the  process  of  exposure  in  the 
air-bath  and  weigh  until  a  constant  weight,  within  reasonable  limits  of  vari- 
ation in  weighing  (0.0003  grm.),  is  obtained. 

Calculate:  (a)  the  percentage  weight  of  water  found  in  the  salt,  (6)  the 
theoretical  percentage  of  water  in  BaCl2  •  2  H2O,  (c)  the  amount  of  water 
which  should  theoretically  be  lost  from  the  actual  weight  of  the  substance 
taken,  and  (d)  the  error  (which  may  be  due  either  to  the  manipulation  or  to 
imperfection  of  the  substance)  in  grams  and  per  cent. 

State  the  results  in  the  following  form: 


Weight  of 
BaCl2  •  2  H2O 
taken. 

Weight  of  water 
found  by  loss. 

Theory  for  water  in 
BaCl2  •  2  H2O 
taken. 

Error. 

Grm. 

Grm. 

Per  cent 

Grm. 

Per  cent 

Grm. 

Per  cent 

*  In  the  case  of  barium  chloride  further  reaction  takes  place  if  the 
temperature  is  too  high  and  the  heating  too  rapid,  hydrochloric  acid  being 
formed  and  eliminated  while  oxygen  replaces  chlorine  in  the  residue. 


36 


QUANTITATIVE  CHEMICAL  ANALYSIS 


Determination  of  Water  by  Loss  on  Heating  with  a 
Flux.  —  Instead  of  dehydrating  crystallized  copper  sulphate, 
CuS04'5H2O,  by  carefully  adjusting  the  conditions  so  that 
the  salt  shall  be  heated  for  a  considerable  time  in  a  current  of 
dry  air  at  a  temperature  gradually  raised  and  not  much  exceed- 
ing 200  degrees,  the  operation  may  be  completed  speedily  and 
safely  if  a  suitable  amount  of  the  easily  fusible  sodium  tung- 
state  *  is  mixed  with  the  copper  sulphate  and  the  mixture  heated 
gradually  to  the  fusing  point.  The  following  expression  may 
serve  as  a  typical  representation  of  this  reaction : 

CuS04  •  5  H2O  +  nNa2W04  =  Na2S04  +  CuO  +  W03 
+  (n-  1)  Na2W04  +  5  H2O. 

Other  sulphates  may  be  similarly  dehydrated  without  losing 
sulphur  trioxide.t 

Experimental  Process,'  Determination  of  Water  in  Copper  Sulphate. — 
Weigh  exactly  in  a  porcelain  crucible  about  0.5  grm.  of  copper  sulphate, 
CuSO4  •  5  H2O,  mix  the  sulphate  well  with  2.5  grm.  of  anhydrous  sodium 
tungstate,  Na2WO4,  and  cover  the  mixture  with  0.5  grm.  of  the  tungstate. 
Weigh  the  crucible  and  mixture,  warm  very  gently  at  first,  and  finally  in- 
crease the  heat  gradually  until  the  fusion  is  complete.  Cool  and  weigh. 
Arrange  the  results  in  the  following  form: 


Copper  sulphate 
taken. 

Loss  on 
heating. 

Theory  for  water  in 
CuSO4  •  5  H2O  taken. 

Error. 

Grm. 

Grm. 

Per  cent 

Grm. 

Per  cent. 

Grm. 

Per  cent 

Determination  of  Water  by  Direct  Weighing.  —  When  sub- 
stances lose  water  upon  application  of  heat  or  upon  fusion  with 
a  flux  and  the  loss  is  accompanied  by  indefinite  change  in  the 

*  To  prepare  the  anhydrous  tungstate,  the  commercial  material, 
Na2WO4,  which  is  ordinarily  contaminated  with  sodium  carbonate,  is 
fused  in  a  platinum  crucible  (or  dish)  and  tungstic  oxide,  WO3  (thoroughly 
ignited  to  insure  the  absence  of  ammonia  and  the  oxidation  of  the  product) 
is  added  until  bubbles  of  carbon  dioxide  cease  to  form. 

t  Kuzirian:  Am.  Jour.  Sci.,  [4]  36,  401  (1913). 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


37 


residue  due  to  atmospheric  action,  or  by  loss  of  other  volatile 
material,  the  water  evolved  may  be  carried  in  a  current  of  air 
to  a  suitable  absorption  apparatus  and  weighed  directly,  pro- 
vided that  the  other  volatile  products  (if  any)  are  held  back 
or  intercepted  before  they  reach  the  water  retainer,  or  else  pass 
through  the  retainer  and  escape  retention  with  the  water. 


FIG.  9. 


FIG.  10. 

Ignition  tube  (Jannasch):  inside  diameter,  1.2  cm. ;  length  (a  to  e),  26 
cm.;  volume  of  ignition-bulb  6,  25  cm.3;  lead  oxide  between  plugs  of  glass 
wool,  c  to  d]  absorption  tube,  /;  protective  tube,  g. 

A  common  method  for  the  direct  determination  of  water  of 
composition  is  to  place  the  substance  in  a  horizontal  combustion 
tube  of  glass  (with  or  without  a .  flux,  and  best  in  a  boat  of 
porcelain  or  platinum),  and  to  heat  the  tube  to  a  suitable  temper- 
ature (in  a  tubulated  air-bath  (Fig.  9),  by  the  direct  flame  of 
burner  or  blast-lamp  (Fig.  10),  or  in  an  electric  furnace)  while 


38  QUANTITATIVE  CHEMICAL  ANALYSIS 

a  current  of  specially  dried  air  passes  over  the  substance  and 
carries  the  evolved  water  to  an  absorption  tube  charged  with 
the  same  kind  of  material  (e.g.,  sulphuric  acid),  which  is  used 
to  dry  the  air  before  it  is  admitted  to  the  tube.* 

The  Determination  of  Carbon  Dioxide 

The  carbon  dioxide  of  carbonates  may  be  determined,  accord- 
ing to  circumstances,  by  loss  on  heating  simply,  by  loss  on  heat- 
ing with  a  flux,  by  loss  of  the  gas  liberated  by  the  action  of  a 
suitable  acid,  or  by  direct  weighing  of  the  gas  evolved  and 
absorbed  in  any  suitable  way. 

Determination  of  Carbon  Dioxide  in  Carbonates  by  Loss 
on  Heating.  —  Excepting  the  carbonates  of  the  alkali  elements, 
barium,  and  strontium,  the  carbon  dioxide  of  fixed  carbonates 
may  be  expelled  by  heating  (ignition)  to  temperature^  easily 
attainable  in  analytical  processes,  and  when  the  residual  sub- 
stance suffers  no  change  of  weight  other  than  that  consequent 
upon  the  loss  of  carbon  dioxide  the  amount  of  the  carbon  dioxide 
may  be  determined  by  finding  the  difference  between  the  weight 
before  ignition  and  the  weight  after  ignition. 

From  the  carbonates  of  magnesium,  zinc,  cadmium,  and  lead, 
the  carbon  dioxide  may  be  expelled  by  simple  ignition  at  the 
temperature  of  the  Bunsen  burner: 

MgCO3  =  MgO  +  C02, 
ZnCO3  =  ZnO  +  C02, 
CdC03  =  CdO  +  C02, 
PbC03  =  PbO  +  CO2. 

In  the  case  of  calcium  carbonate  the  process  of  decomposition 
in  the  reversible  reaction 

CaCO3<=±CaO  +  C02 

is  practically  complete  only  at  the  high  heat  of  the  blast-lamp 
and  when  the  resulting  carbon  dioxide  is  removed. 

*  For  details  in  respect  to  apparatus  used  in  the  determination  of  water 
in  minerals  and  rocks,  see  Hillebrand:  Bulletin  No.  422  of  the  U.  S.  Geo- 
logical Survey. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


39 


Experimental  Process:  Determination  of  Carbon  Dioxide  in  Calcite  by 
Ignition.  —  Weigh  out  exactly  in  a  weighed  platinum  crucible  about 
0.5  grm.  of  powdered  calcite.  Heat  the  crucible  and  contents  for  five 
minutes  well  above  the  Bunsen  flame  turned  low.  Cool  and  weigh  again. 
The  loss  of  weight  should  represent  hygroscopic  moisture.  Take  the 
difference  between  the  weight  of  the  crucible  and  the  weight  obtained  after 
the  gentle  ignition  as  the  weight  of  dry  calcite.  Now  ignite  fully  for  half 
an  hour  in  a  strong  Bunsen  flame  (keeping  the  crucible  above  the  blue 
cone  of  the  flame,  to  avoid  unfavorable  action  of  the  unburned  gas  upon 
the  platinum)  and  finally  over  the  blast-lamp  for  five  minutes.  Cool  in 
the  desiccator  and  weigh.  Repeat  the  five-minute  ignition  over  the  blast- 
lamp  as  many  times  as  may  be  necessary  to  secure  a  practically  constant, 
weight  (i.e.,  a  weight  which  does  not  differ  from  the  preceding  weight  by 
more  than  0.0003  grm.).  Note  the  loss  found  in  igniting  the  dried  calcite 
to  constant  weight,  and  calculate  the  percentage  loss.  Compare  the  re- 
sults obtained  with  those  which  should  have  been  obtained  if  the  calcite 
were  pure  CaC03  and  the  analytical  process  perfect.  State  the  results  in 
the  following  form: 


Weight  of 
calcite 
taken. 

Loss  on 
drying. 

Loss  on  ignition. 

Theory  for  loss  in 
ignition  of  pure 
dried  calcite. 

Error. 

Grm. 

Grm. 

Grm. 

Per  cent. 

Grm. 

Per  cent. 

Grm. 

Per  cent. 

Determination  of  Carbon  Dioxide  in  Carbonates  by  Fusion 
with  a  Flux.  —  Some  carbonates  —  strontium  carbonate, 
barium  carbonate,  and  the  carbonates  of  the  alkali  elements  — 
do  not  yield  their  content  of  carbon  dioxide  in  a  simple  ignition 
under  conditions  ordinarily  available.  Such  carbonates,  as 
well  as  the  less  refractory  carbonates,  may  be  satisfactorily 
decomposed  by  the  aid  of  a  suitable  flux  to  combine  with  the 
residual  oxide  in  a  non-reversible  reaction  and  thus  to  assist  in 
the  expulsion  of  volatile  carbon  dioxide.  Sodium  paratung- 
state,  Nai0Wi204i,  makes  a  very  good  and  easily  fusible  flux  for 
the  treatment  of  the  refractory  carbonates.*  Approximately 

*  This  material  may  be  prepared  by  heating  to  fusion  over  the  blast- 
lamp  a  known  weight  of  sodium  tungstate,  NaaWC^  •  2  H2O,  adding  an 
equal  weight  of  previously  ignited  tungsten  trioxide,  WO3,  heating  to  the 
clear  fusion,  cooling,  and  grinding  in  a  mortar. 


40  QUANTITATIVE  CHEMICAL  ANALYSIS 

half  the  weight  of  the  paratungstate  is  tungsten  trioxide  (equiv- 
alent weight,  232)  and  this  should  be  capable  of  expelling  car- 
bon dioxide  (molecular  weight,  44)  to  the  amount  of  a  fifth  of 
its  own  weight.  The  weight  of  paratungstate  employed  in  the 
analytical  determination  should,  therefore,  be  somewhat  in  ex- 
cess of  ten  times  the  weight  of  carbon  dioxide  to  be  expelled. 

Experimental  Process:  Carbon  Dioxide  in  Calcite.  —  Put  in  a  weighed 
crucible  (platinum  or  porcelain)  about  0.5  grm.  of  powdered  calcite. 
Heat  the  crucible  and  contents  for  five  minutes  well  above  the  Bunsen 
flame  turned  low.  Cool  and  weigh.  Take  the  difference  between  the 
weight  of  the  crucible  alone  and  the  crucible  with  contents  after  drying  as 
the  weight  of  the  dried  calcite.  Now  add  about  2.5  grm.  of  previously 
dried  sodium  paratungstate,  NaioWuO*!.  Stir  carefully  with  a  platinum 
wire  taking  care  to  avoid  mechanical  loss,  and  weigh  again.  Heat  the 
crucible  over  a  Bunsen  flame,  first  at  very  low  heat  and  then  to  fusion  of 
the  mixture  for  five  minutes.  Cool  in  a  desiccator  and  weigh.  « Repeat 
the  ignition  to  constancy  of  weight,  which,  however,  is  usually  obtained 
in  the  first  fusion.  Note  the  weight  of  carbon  dioxide  lost  from  the  dry 
calcite,  calculate  the  percentage  loss,  and  compare  the  results  obtained 
with  those  which  should  have  been  obtained  if  the  material  were 
pure  and  the  process  perfect.  State  the  results  in  the  form  given 
on  p.  39. 

Determination  of  Carbon  Dioxide  Evolved  by  the 
Action  of  an  Acid  upon  Carbonates.  —  The  carbon 
dioxide  of  a  carbonate  may  be  liberated  by  the  action 
of  a  suitable  acid,  and  in  absence  of  other  volatile  sub- 
stances it  may  be  determined  by  the  loss  in  weight  of 
a  suitably  designed  apparatus. 

Various  forms  of  apparatus,  some  of  them  very 
elaborate,  have  been  devised  to  determine,  by  loss  of 
weight,  carbon  dioxide  and  other  volatile  products  of 
FIG.  11.  reaction.  A  convenient  and  effective  device  for  such 
purposes  is  the  Kreider  apparatus,  made  very  simply 
from  three  test-tubes,  modified  and  fitted  as  shown  in  the 
accompanying  figure.  The  test-tube  A  serves  as  the  reaction 
chamber;  B,  the  drying  chamber,  is  perforated  with  a  hole 
about  1  cm.  in  diameter  and  fits  tightly  within  A',  and  the 
reagent  chamber  C,  so  selected  that  it  fits  loosely  within  B, 
is  drawn  out  to  a  small  capillary  tube.  To  C  is  fitted  a  one- 
holed  stopper,  through  which  passes  a  short  glass  tube  which  is 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  41 

provided  with  a  rubber  cap  and  plug  or  with  a  stop-cock. 
Upon  removing  the  plug  or  opening  the  stop-cock  and  applying 
suction  to  the  short  tube,  the  reagent  employed  to  liberate  the 
volatile  product  may  be  drawn  up  through  this  capillary  until 
C  is  sufficiently  filled  with  acid  and  the  lower  end  of  the  capil- 
lary with  air. 

When  the  apparatus  is  to  be  used,  the  capillary  of  C,  which 
has  been  fitted  as  described,  is  pushed  through  the  hole  of  B, 
packed  loosely  with  cotton;  B  is  filled  to  the  depth  of  from  6  cm. 
to  8  cm.  (about  two-thirds  of  its  contents),  with  granular  cal- 
cium chloride;  and  B  and  C  are  adjusted  as  shown.  The 
substance  to  be  analyzed  is  weighed  and  put  in  A.  The  reagent 
to  be  employed  (10  cm.3  to  15  cm.3)  is  drawn  into  C,  and  held 
there  in  the  manner  described  above.  The  test-tube  A  is 
slipped  over  B,  and  a  ring  of  paraffin  is  melted  into  the  mouth 
of  A,  about  B,  by  means  of  a  hot  iron  or  wire,  to  seal  the  joint 
securely.  The  apparatus  is  wiped,  hung  in  the  balance,  and, 
after  a  suitable  interval,  weighed.  A  very  convenient  way  to 
attach  the  paraffin  is  to  melt  it  between  A  and  another  tube, 
which  fits  A  just  as  B  does,  and  which  may  be  removed  by  a 
turning  motion,  leaving  the  ring  into  which  B  will  fit.  Very 
little  heating  is  then  required  to  make  a  tight  joint.  If  care  be 
used  in  taking  apart  A  and  B,  at  the  close  of  an  experiment, 
such  a  ring  of  paraffin  remains  in  place  and  may  be  used  many 
times  without  replacement,  being  remelted  by  a  touch  of  the 
hot  wire  before  every  new  experiment. 

Experimental  Process:  Carbon  Dioxide  in  Caldte.  —  Weigh  out  exactly 
upon  a  balanced  watch-glass  about  0.5  grm.  of  powdered  calcite.  Place 
in  the  mouth  of  the  upright  reaction  chamber  A  of  the  Kreider  apparatus, 
just  described,  a  funnel  having  the  stem  cut  off  close  to  the  cone.  Trans- 
fer nearly  all  of  the  weighed  calcite  to  the  funnel  by  inclining  the  watch- 
glass  and  tapping  it  on  the  edge  with  the  quill  of  a  camel's-hair  brush  and 
complete  the  transfer  by  touching  the  glass  lightly  with  the  tip  of  the 
brush,  taking  care  to  remove  to  the  funnel  any  adherent  particles  of  the 
powder  by  bringing  the  brush  lightly  upon  the  edge  of  the  watch-glass. 

Fill  the  reagent  chamber  C,  in  the  manner  described,  with  10  cm.3  to  15 
cm.3  of  dilute  hydrochloric  acid  (1  :  3).  Put  the  apparatus  together,  seal- 
ing the  paraffin  joint,  and  hang  it  upon  the  balance.  Weigh  after  an  inter- 
val of  twenty  minutes. 

Remove  the  cap  from  the  small  tube  in  the  stopper  of  the  reagent 
chamber  C  or  open  the  stop-cock  and  allow  the  acid  to  run  from  C  into  A . 
The  carbon  dioxide  liberated  is  forced  upward  through  the  drying  chamber 
B  and  escapes  through  the  annular  space  between  B  and  C.  When  the 
visible  action  has  ceased,  force  a  gentle  current  of  air,  dried  by  passing  it 


42 


QUANTITATIVE  CHEMICAL  ANALYSIS 


through  calcium  chloride,  through  C  to  remove  the  carbon  dioxide  from 
the  apparatus.  Replace  the  cap,  hang  the  apparatus  on  the  balance,  and 
weigh  after  twenty  minutes.  The  loss  of  weight  represents  the  carbon 
•dioxide. 

Calculate  the  weight  of  carbon  dioxide  in  the  weight  of  calcite  taken, 
and  the  percentage  of  carbon  dioxide,  and  compare  the  results  with  those 
which  should  be  obtained  upon  the  theory  that  the  calcite  is  pure  and  the 
analytical  process  free  from  error.  State  the  details  of  the  analysis  in  the 
following  form: 


Weight  of  calcite 
taken. 

Loss  in  weight  of 
apparatus. 

Theory  for  loss  of  CO2. 

Error. 

Grin. 

Grin. 

Grm. 

Per  cent. 

Grm. 

Per  cent. 

4 

Determination  of  Carbon  Dioxide  by  Direct  Weighing.  - 

Carbon  dioxide  evolved  from  carbonates  by  heating  either  with 
or  without  a  flux,  or  liberated  by  the  action  of  an  acid,,  may 
be  dried  by  passing  it  through  some  drying  material  which 
is  incapable  of  retaining  it  and  may  then  be  absorbed  in 
some  suitable  retainer  like  potassium  hydroxide  or  soda-lime. 
Usually  when  carbon  dioxide  is  to  be  weighed  directly  it  is 
evolved  by  the  action  of  an  acid. 

A  convenient  apparatus  for  the  evolution  of  carbon  dioxide 
from  carbonates  by  the  action  of  an  acid  and  the  absorption  of 
that  gas  in  a  suitable  retainer  for  weighing  is  shown  in  Fig.  12.* 
This  apparatus  consists  of  an  evolution  flask,  a  cooler  to  return 
condensed  steam  to  the  flask,  a  U-tube  A  filled  with  calcium 
chloride,  a  double  U-tube  B  charged  in  the  central  section  with 
anhydrous  copper  sulphate  (to  intercept  traces  of  hydrochloric 
acid  which  may  escape  previous  condensation)  with  calcium 
chloride  to  dry  the  gas,  two  absorption  tubes  C  and  D  filled  with 
soda-lime  (followed  by  calcium  chloride)  to  retain  the  carbonic 
acid,  and  gas  observation  bulbs  E  filled  with  sulphuric  acid  to 
show  the  rate  of  gas  flow.  Through  the  rubber  stopper  of  the 
flask  is  passed  the  tube  of  a  separating  funnel,  the  end  of  the  con- 
densing tube,  and  a  tube  (which  connects  the  interior  of  the  flask 

*  Hillebrand:  Bulletin  422  of  the  U.  S.  Geological  Survey,  p.  179. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


43 


with  the  system  of  bottles  and  the  purifying  tower)  to  supply 
under  slight  pressure,  air  free  from  carbon  dioxide. 

In  carrying  out  the  determination  of  carbon  dioxide,  the 
powdered  carbonate  is  introduced  into  the  flask  and  covered 
with  hot,  boiled  water.  A  current  of  air,  freed  from  carbon 
dioxide  by  passing  through  the  purifier  charged  with  soda-lime, 


FIG.  12. 


is  forced  through  that  part  of  the  system  which  precedes  the 
weighed  absorption  bulbs  and  tube,  to  displace  the  unpurified 
air.  The  stop-cock  in  the  separating  funnel  is  closed  and  the 
funnel  is  charged  with  hydrochloric  acid  (1  :  1).  The  connec- 
tion is  made  between  the  weighed  absorption  apparatus  and 
the  rest  of  the  system.  The  liquid  in  the  flask  is  brought  to  the 
boiling  point  and  the  acid  is  allowed  to  flow  into  the  flask  and 
purified  air  is  sent  slowly  through  the  boiling  liquid  into  which 
the  end  of  the  aspirating  tube  and  the  end  of  the  funnel  tube  dip. 
At  the  end  of  half  an  hour  the  burner  is  removed  from  under  the 
flask  and  the  apparatus  is  permitted  to  cool  for  another  half 
hour  in  the  current  of  purified  air.  The  difference  between  the 


44  QUANTITATIVE  CHEMICAL  ANALYSIS 

weights  of  the  potash  bulbs,  and  the  accompanying  drying  tube 
before  and  after  the  operation  is  the  weight  of  carbon  dioxide 
absorbed. 

Determination  of  Carbon  as  Carbon  Dioxide  and  Hydrogen 
as  Water :  by  the  Combustion  Process.  —  Organic  compounds 
containing  carbon  and  hydrogen  may  be  decomposed  by 
ignition  with  copper  oxide  in  oxygen.  The  hydrogen  forms 
water  which  is  absorbed  by  granulated  calcium  chloride.  The 
carbon  forms  carbon  dioxide  which  passes  through  the  cal- 
cium chloride  (free  from  oxide)  and  is  absorbed  by  an  alkali 
hydroxide.  In  estimating  the  carbon  dioxide  formed  in  the 
process  of  combustion  the  apparatus  is  arranged  so  that  com- 
bustion shall  be  complete,  the  water  taken  up  by  weighed 
calcium  chloride,  and  the  residual  carbon  dioxide  passed  on  to 
an  absorption  tube  containing  a  weighed  amount  of  potassium 
hydroxide.* 

The  Determination  of  Boric  Acid 

Free  boric  acid  reacts  with  an  excess  of  methyl  alcohol  to  form 
methyl  orthoborate  which  boils  at  65  degrees. 

3  CH3OH  +  H303B  <=±  (CH3)03B  +  3  H20. 
When  methyl  borate  meets  calcium  hydroxide  moistened  with 
water  calcium  borate  is  formed  and  methyl  alcohol  is  set  free. 
rcCa02H2  +  2  (CH3)303B  +  2  H20  =  Ca02(BO)2 

+  6  CH3OH  +  (n  -  l)Ca02H2. 

Similarly  when  methyl  borate  and  sodium  tungstate  react  in 
solution  a  sodium  borate  and  methyl  alcohol  result  in  the  sense 
of  the  following  representative  equation: 

nNa2W04  +  2  (CH3)303B  +  3  H20  =  2  NaOBO 

+  6  CH3OH+  (n  -  l)Na2W04  +  W03.f 

Upon  treating  boric  acid  or  borates  of  the  alkali  elements  with 
absolute  methyl  alcohol  (free  from  acetone)  and  acetic  acid  the 

*  For  a  description  of  apparatus  and  procedures  in  the  determination 
of  carbon,  hydrogen,  and  nitrogen  by  the  combustion  process,  see  Tread- 
well-Hall:  Analytical  Chemistry,  3rd  ed.,  Vol.  II,  pp.  414  to  426. 

f  In  combination. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


45 


volatile  methyl  borate  is  formed  which  upon  treatment  with 
calcium  hydroxide  formed  from  a  weighed  amount  of  calcium 
oxide,  or  with  a  known  weight  of  sodium  tungstate,  is  decom- 
posed with  the  elimination  of  alcohol  and  the  formation  of  a 
residue,  the  weight  of  which  after  ignition  differs  from  that  of 
the  original  calcium  oxide,  or  sodium  tungstate,  by  the  amount 
of  the  weight  of  boron 
trioxide  absorbed.  These 
reactions  have  been  made 
the  basis  of  analytical  proc- 
esses for  the  determination 
of  boron  trioxide  in  boric 
acid  or  borates.* 

The  Apparatus.  -  -  The 
distillation  apparatus,  which 
is  shown  in  the  accompany- 
ing cut,  Fig.  13,  consists 
essentially  of  a  retort,  con- 
denser, and  a  paraffin  bath 
for  heating.  The  condenser 
is  set  vertically,  to  facilitate 
changing  the  level  of  the 
retort  within  the  bath,  and 
to  secure  at  the  same  time 
continual  and  thorough 
washing  of  the  tube  by  its 
own  condensations.  The  retort,  somewhat 


FIG.  13. 


like  the  well- 
known  drying  tube  of  Liebig  in  general  shape,  is  easily 
made  from  a  pipette  by  bending  the  tube  at  one  end  to 
a  right  angle,  at  the  other  to  a  goose-neck,  as  shown.  To  the 
former  is  fitted  or  sealed  a  glass  funnel-tube  provided  with  a 
stop-cock.  The  end  of  the  goose-neck  passes  tightly  through  a 
rubber  stopper  in  the  upper  end  of  the  condensing  tube.  A 
small  Erlenmeyer  flask  which  is  joined  to  the  condenser,  by 
means  of  a  thistle-tube  and  a  rubber  stopper  grooved  to  per- 
mit the  free  passage  of  air,  receives  the  distillate.  In  carrying 

*  Rosenbladt:  Zeit.  Anal.  Chem.,  27,  18. 
Gooch:  Proc.  Am.  Acad.,  22,  167  (1886);  Am.    Chem.  Jour.,  9,  23 

(1887). 
Gooch  and  Jones:  Am.  Jour.  Sci.,  [4],  7,  34. 


46  QUANTITATIVE  CHEMICAL  ANALYSIS 

out  a  distillation,  the  liquid  to  be  distilled  is  introduced  into 
the  retort  by  the  funnel-tube,  the  glass  cock  is  closed,  the  water 
started  through  the  condenser,  and  the  retort  lowered  into  the 
hot  paraffin,  care  being  taken  to  begin  the  operation  with  the 
retort  not  more  than  half  full  and  so  inclined  that  only  the  rear 
dips  below  the  surface  of  the  bath.  If  the  precaution  to  heat 
the  retort  at  the  start  in  this  manner  be  overlooked,  it  may 
sometimes  happen  that  the  sudden  and  violent  expulsion  of  air 
through  the  liquid  will  carry  portions  of  it  bodily  into  the  goose- 
neck, and  even  into  the  condenser.  With  this  point  considered, 
the  remainder  of  the  operation  presents  no  difficulty,  and  re- 
quires little  care. 

The  size  of  the  retort  may  be  suited,  of  course,  to  the  particu- 
lar case  in  hand,  but  for  most  purposes  a  200-cm.3  pipette  makes 
a  retort  of  convenient  dimensions,  neither  too  large  for  the  dis- 
tillation of  small  charges  nor  too  small  to  permit  the  treatment 
of  100  cm.3  of  liquid  comfortably.  The  tube  of  the  goose-neck 
should  be  wide  enough  to  prevent  the  formation  of  bubbles  in 
it;  0.7  cm.  is  a  good  measure  for  the  interior  diameter.  It  is  of 
advantage  to  heat  the  bath  to  a  point  considerably  above  the 
temperature  at  which  the  liquid  which  is  to  be  distilled  boils,  — 
something  between  130°  and  140°  does  very  well  for  water, 
and  is  not  too  high  for  methyl  alcohol,  —  and  under  such  cir- 
cumstances, and  when  the  retort  is  entirely  submerged,  it  often 
happens  that  evaporation  takes  place  with  extreme  rapidity 
from  the  surface  of  the  liquid  in  perfect  quiet  without  actual 
boiling. 

The  number  of  distillations  necessary  depends,  of  course, 
upon  the  amount  of  boric  acid  treated.  To  remove  0.2  grm. 
of  boric  anhydride  completely  to  the  distillate,  six  charges  of 
methyl  alcohol  of  10  cm.3  should  be  ample. 

In  treating  borates  soluble  in  water,  the  solution,  made  barely 
acid  to  litmus  (added  to  the  solution)  by  means  of  hydro- 
chloric acid,  then  just  alkaline  by  sodium  hydroxide,  and  again 
faintly  acid  by  acetic  w  acid,  is  put  in  the  retort  through  the 
funnel  tube  which  is  washed  down  by  two  or  three  cubic  cen- 
timeters of  water  used  in  three  portions,  successively. 

Use  of  Calcium  Oxide  as  the  Retainer.  —  When  calcium  oxide 
is  to  be  used  to  retain  the  boric  acid  in  the  distillate  it  is  first 
ignited  in  the  crucible  in  which  the  evaporation  of  the  distillate 
is  to  be  made  subsequently,  and  then  transferred  to  the  receiv- 
ing flask  which  is  attached  to  the  condenser.  To  prevent  the 
caking  of  the  lime  by  the  action  of  the  alcohol,  it  is  slaked 
with  a  little  water  before  the  distillation. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  47 

The  retort  is  connected  with  the  condenser  and  submerged 
in  the  paraffin  at  a  temperature  not  exceeding  140  degrees. 
When  all  of  the  liquid  has  passed  to  the  distillate  the  retort  is 
raised  from  the  bath,  allowed  to  cool,  and  the  dry  residue  is 
drenched  with  10  cm.3  of  methyl  alcohol  (free  from  acetone), 
the  retort  is  again  submerged  in  the  paraffin,  and  the  methyl 
alcohol  is  distilled  to  the  receiver.  After  three  such  treat- 
ments with  methyl  alcohol,  the  residue  in  the  retort  is  mois- 
tened with  a  very  little  water  (2  cm.3)  and  the  blue  color  of  the 
litmus  (due  to  the  hydrolytic  decomposition  of  the  acetate  pres- 
ent) is  changed  to  red  by  the  addition  of  a  drop  or  two  of  acetic 
acid.  Thereafter  the  treatment  with  methyl  alcohol,  applied 
in  10-cm.3  portions,  is  repeated  three  times,  and  the  receiver 
is  then  detached  from  the  condenser,  stoppered,  shaken  thor- 
oughly, and  allowed  to  stand  fifteen  minutes.  The  contents  of 
the  receiver  are  poured  and  washed  out  as  completely  as  may 
be  into  a  large  platinum  dish  and  the  alcoholic  liquid  is  slowly 
evaporated  to  dryness  at  a  temperature  below  that  at  which  it 
will  boil.  The  residue  is  transferred  by  means  of  a  little  water 
to  the  crucible  in  which  the  calcium  oxide  was  originally  ignited. 
Any  material  adhering  to  the  walls  of  the  receiver  and  of  the 
dish  is  dissolved  in  a  very  few  drops  of  dilute  nitric  acid  and 
added  to  the  contents  of  the  crucible  which  are  evaporated  to 
dryness  on  the  water-bath,  ignited,  gently  at  first  and  finally  to 
constant  weight. 

Use  of  Sodium  Tung  state  as  the  Retainer.  —  Sodium  tungstate, 
fused  with  a  slight  excess  of  tungstic  acid  over  that  contained 
in  the  normal  tungstate  (to  insure  its  freedom  from  carbonate), 
may  be  used  with  good  results  as  an  absorbent  for  boric  acid. 
This  substance  is  definite  in  weight,  not  hygroscopic,  soluble 
in  water,  and  recoverable  in  its  original  weight  after  evapo- 
ration and  ignition. 

In  this  procedure  use  is  made  of  the  apparatus  described 
above  and  shown  in  Fig.  13,  modified  in  the  single  point  that 
the  Erlenmeyer  flask  used  as  a  receiver  is  fitted  tightly  to  the 
condenser  and  provided  with  a  water  trap.  Special  care 
should  be  taken  to  give  the  tungstate  ample  time  for  contact 
with  the  distillate  before  exposing  the  latter  to  atmospheric 
evaporation.  The  distillate  is  received,  therefore,  in  a  dilute 
solution  of  sodium  tungstate  which  is  placed  in  the  receiver, 
cooled  by  ice,  and  trapped  with  water.  The  mixture  is  well 
stirred,  allowed  to  stand  half  an  hour  after  the  distillation  is 
complete,  evaporated  to  small  volume  in  a  large  dish,  and 
transferred  to  the  crucible  in  which  the  tungstate  was  originally 


48  QUANTITATIVE  CHEMICAL  ANALYSIS 

weighed.  After  thorough  drying  the  residue  is  ignited  to 
fusion  and  weighed.  When  acetic  acid  is  employed  in  the  re- 
tort, care  must  be  taken  in  the  ignition  to  expose  the  fused 
mass  freely  to  the  air  (by  causing  it  to  flow  upon  the  sides  of  the 
crucible)  until  the  color  of  the  cooled  tungstate  is  white,  in 
order  that  the  reducing  effect  of  the  acetate  may  be  eliminated. 
In  determining  boron  trioxide  in  insoluble  substances,  such 
as  silicates,  the  material  is  fused  with  four  times  its  weight  of 
sodium  carbonate  and  the  melt  is  extracted  with  water.  The 
solution  is  evaporated  to  small  volume,  hydrochloric  acid  is 
added  to  faint  acidity,  a  drop  or  two  of  a  solution  of  sodium 
hydroxide  to  alkalinity,  and  acetic  acid  to  acid  reaction  to 
litmus.  The  solution  thus  prepared  may  be  introduced  into 
the  retort  and  treated  in  the  manner  described  above  for  the 
determination  of  boron  trioxide  in  soluble  borates. 

The  Determination  of  the  Fixed  Product  of  Action 
of  a  Volatile  Reagent 

The  determination  of  the  fixed  product  of  the  action  of  a 
volatile  reagent  upon  a  substance  may  be  the  object  of  analysis, 
and  it  is  frequently  desirable  to  make  use  of  a  reagent,  by  itself 
volatile,  which  will  enter  into  definite  reaction  with  the  sub- 
stance to  form  a  definite  and  weighable  product. 

Determination  of  Magnesium  and  Manganese  as  Anhy- 
drous Sulphates.  —  Examples  of  such  processes  are  the  deter- 
mination of  magnesium  and  the  determination  of  manganese, 
each  in  the  form  of  the  anhydrous  sulphate  derived  by  the 
action  of  sulphuric  acid  in  excess  upon  salts  of  these  elements, 
the  excess  of  the  sulphuric  acid,  as  well  as  the  volatile  product 
liberated  in  the  action,  being  expelled. 

Experimental  Process:  Conversion  of  Magnesium  Chloride  to  Magnesium 
Sulphate.  —  Draw  from  a  burette  into  a  roomy  crucible  a  definite  volume 
(50  cm.3)  of  a  solution  of  magnesium  chloride  (about  N/10  in  respect  to 
the  anhydrous  salt),  add  sulphuric  acid  (about  1  cm.3  of  the  concentrated 
acid,  or  its  equivalent),  evaporate  the  solution  on  the  water-bath,  suspend 
the  crucible  in  another  crucible  (to  be  used  as  a  radiator)  with  the  walls 
about  1  cm.  apart,  heat  the  outer  crucible  with  gradually  rising  tem- 
perature until  the  free  sulphuric  acid  has  been  removed  and  then  to 
incipient  redness,  and  weigh  the  residue  as  magnesium  sulphate,  MgSO4. 

Calculate  and  state  the  magnesium  content  per  cm.3  of  the  solution. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


49 


Volume  of  solution 
taken. 

Weight  of  MgSO4 
found. 

Weight  of  magnesium 
calculated. 

Magnesium  content 
per  cm.3 

Cm.3 

Grm. 

Grm. 

Grm. 

Experimental  Process:  Determination  of  Manganese  in  the  Hydrous  Sul- 
phate. —  Weigh  out  exactly  in  a  crucible  about  1  grm.  of  hydrous  man- 
ganous  sulphate  (MnS04  •  nH2O).  Place  the  crucible  upon  a  triangle 
within  a  larger  porcelain  crucible  in  such  manner  that  the  walls  of  the 
inner  crucible  shall  be  distant  from  that  of  the  outer  crucible  by  about 
1  cm.  Heat  the  outer  crucible  gradually  until  the  bottom  is  at  red  heat. 
Continue  the  heating  (at  a  temperature  of  400  to  450  degrees)  weighing 
at  intervals,  until  the  weight  is  constant.  State  the  results  in  the  form 
given  below,  finding  by  calculation  the  numerical  symbol  which  best  repre- 
sents the  indefinite  symbol  MnS04  •  nH20. 


Weight  of 
MnS04  •  nH2O 
taken. 

Weight  of  MnSO4 
found. 

Weight  of  H2O  in 
MnSO4  •  nH2O  found. 

Numerical  symbol 
of  MnSO4  •  nH2O. 

Grm. 

Grm. 

Grm. 

ELECTROLYTIC   PROCESSES 

Many  valuable  methods  of  electro-analysis,  chiefly  determina- 
tions of  the  metals,  are  based  upon  the  decomposition  of  electrolytes 
in  solution.  Solutions  of  such  substances  carry  the  electric 
current  and  the  dissolved  electrolytes  are  decomposed  at  the 
electrodes,  with  the  separation  of  an  anion  at  the  anode  (posi- 
tive pole)  and  a  cation  at  the  cathode  (negative  pole).  The 
current  necessary  for  the  electrolysis  may  be  derived  from 
primary  cells,  storage  cells,  or  the  electric-lighting  circuit,  as 
may  be  convenient. 

The  Conditions  of  Electrolysis 

According  to  the  prevailing  theory  of  electrolysis,  all  the  ions 
of  a  solution,  of  whatever  nature,  are  acted  upon  by  the  electric 


50  QUANTITATIVE  CHEMICAL  ANALYSIS 

force  and  all  carry  the  current  by  moving  through  the  solu- 
tion. If  more  than  one  kind  of  ion  of  the  same  sign  is  present 
that  kind  which  has  the  lower  decomposition  voltage  is  first 
deposited  at  the  corresponding  electrode. 

In  the  decomposition  of  a  metallic  salt  in  a  water  solution 
from  which  the  metal  can  be  deposited  by  action  of  the  current, 
the  metal  ion  is  deposited  at  the  cathode,  while  the  acid  ion  is 
liberated  at  the  anode.  If  the  concentration  of  the  electrolyte 
near  the  cathode  falls  to  a  point  at  which  the  solution  is  deficient 
in  metal  ions  to  deliver  the  current  to  the  cathode,  hydrogen 
ions  of  the  water  or  acid  may  be  discharged  simultaneously 
with  the  metal  and  the  continuity  of  the  metallic  deposit 
affected.  In  order  that  smooth  and  adherent  deposits  may  be 
obtained  it  is  generally  necessary  to  regulate  the  rate  of  dep- 
osition of  the  metal  by  keeping  the  current  within  define^  limits 
of  surface-density  —  that  is,  within  certain  limits  which  may  be 
expressed  in  terms  of  amperes  per  100  cm.2  of  electrode  surface, 
or  normal  density  (JVDioo). 

amperes  X  100 


NDm  = 


cm.2  of  electrode  surface 


With  stationary  electrodes  the  surface  density  must  be  low  and 
the  deposition  of  the  metal  consequently  slow,  in  order  that 
the  deposits  may  be  smooth  and  adherent. 

In  certain  electrolytic  processes  which  involve  separations  of 
cations  from  one  another,  or  anions  from  one  another,  careful 
attention  must  also  be  given  to  the  current  pressure  (voltage), 
since,  at  normal  concentrations,  every  electrolyte  has  a  definite 
decomposition  pressure. 

Stationary  Electrodes.  —  A  platinum  dish  of  200  cm.3 
capacity  is  very  serviceable,  both  as  the  container  and  as  the 
electrode  to  receive  the  deposit,  in  processes  conducted  with 
stationary  electrodes;  and  for  the  second  electrode,  generally 
the  anode,  a  stiff  platinum  wire  bent  at  one  end  to  a  flat  spiral 
is  convenient  (Fig.  I4ab).  Many  forms  of  cathodes  and 
anodes  have  been  described  for  special  purposes,  or  as  pos- 
sessing special  advantages  for  general  purposes.  Electrodes  of 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


51 


platinum  gauze  possess  the  advantage  of  being  permeable  to 
the  solution  and  of  receiving  the  deposit  on  the  entire  surface 
of  each  separate  wire,  so  that  the  electrode  surface  is  very  large 
in  comparison  with  that  of  a  dish  electrode 
or  foil  electrode  of  similar  weight. 

Experimental  Process:  Copper  in  Copper  Sulphate.  — 
Dissolve  in  100  cm.3  of  water,  in  a  platinum  dish, 
about  0.5  grm.  of  copper  sulphate,  CuSO4  •  5  H2O, 
weighed  exactly,  and  add  5  cm.3  of  concentrated  nitric 
acid.  Set  the  dish,  which  is  to  be  used  as  the  cathode, 
upon  a  metal  triangle  connected  with  the  negative 
pole  of  a  storage  battery.  Connect  the  platinum 
anode  (e.g.,  the  spiral  of  Fig.  146)  with  the  positive 
pole  of  the  battery  and  immerse  it  in  the  solution. 
Cover  the  dish  with  a  split  watch-glass.  The  battery 
should  be  arranged  to  deliver  the  current  at  a 
potential  of  2  volts  at  the  electrodes.  Under  the 
conditions  described,  the  electrolysis  may  be  com- 


FIG.  14. 


pleted  in  the  cold  over  night,  or  in  a  few  hours  if  the  solution  is  kept 
at  70  to  80  degrees  by  means  of  a  heated  disk  of  wire  gauze  or  asbestos 
placed  about  a  centimeter  below  the  dish.  Determine  the  end  of  the  de- 
composition by  diluting  the  solution  slightly  with  water  and  noting  whether 
copper  is  deposited  in  the  course  of  ten  minutes  upon  the  newly  covered 
cathode  surface.  To  avoid  the  action  of  the  strong  free  acid  upon  the 
deposited  copper,  either  siphon  off  the  liquid  while  introducing  distilled 
water,  until  the  current  ceases,  or  else -add  a  solution  of  ammonium  acetate, 
made  by  neutralizing  with  acetic  acid  at  least  as  much  ammonia  as  is  equiv- 
alent to  the  nitric  acid  taken  and  the  sulphuric  acid  set  free.  Remove 
the  dish,  wash  the  deposit  with  water  and  then  rinse  with  alcohol.  Dry 
the  dish,  with  the  deposit,  by  heating  it  in  an  air-bath  at  a  temperature 
below  100  degrees,  or  by  passing  it  back  and  forth  over  and  well  above  a 


52 


QUANTITATIVE  CHEMICAL  ANALYSIS 


flame  until  the   film   of  moisture  vanishes.     Cool   in   a  desiccator  and 
weigh. 

Compare  the  amount  of  copper  found  with  that  which  should  be  present 
theoretically  in  tne  amount  of  copper  sulphate  taken.  State  the  results 
in -the  following  form: 


Copper 
sulphate  taken. 

Copper  found. 

Theory  for  copper  in  copper 
sulphate. 

Error. 

Grm. 

Grm. 

Grm. 

Per  cent. 

Grm. 

Per  cent. 

The    Rotating   Electrode.  —  If   the   liquid    be    stirred    by 
rotation  of  either  electrode  (or  otherwise)  the  available  range  of 

current  density  may  be  much  in- 
creased. 

Satisfactory  and  very  rapid 
depositions  of  metals  may  be 
made  without  the  use  of  special 
or  expensive  electrodes  by  utiliz- 
ing the  common  platinum  crucible 
as  a  rotating  cathode.  An  ordi- 
nary (20-cm.3)  platinum  crucible 
may  be  rotated  at  a  speed  as  high 
as  800  revolutions  per  minute  by 
means  of  a  small,  inexpensive 
electric  motor  fastened  so  that 
its  shaft  is  vertical  (Fig.  15). 
Upon  this  shaft  the  crucible  is 
fixed  by  pressing  it  over  a  rubber 
stopper  bored  centrally  and  fitted 
tightly  on  the  end  of  the  shaft 
(Fig.  15A).  To  secure  electrical 
connection  between  crucible  and  shaft,  a  narrow  strip  of  sheet 
platinum  is  soldered  to  the  shaft  and  then  bent  upward  along 
the  sides  of  the  stopper,  thus  putting  the  shaft  in  contact  with 
the  inside  of  the  crucible  when  the  latter  is  pressed  over  the 


FIG.  15. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  53 

stopper.  As  a  matter  of  convenience  in  removing  the  crucible, 
the  shaft  is  made  in  two  parts  which  are  joined,  with  care  to 
make  a  good  contact  between  the  two  pieces  of  shafting,  by  a 
rubber  connector  of  sufficient  thickness  to  prevent  the  crucible 
from  wabbling  when  rotated. 

The  solution  to  be  electrolyzed  is  placed  in  a  beaker  upon  a 
small  adjustable  stand,  so  that  the  crucible  may  be  dipped 
into  the  liquid  to  any  desired  depth.  A  platinum  plate  is  em- 
ployed as  an  anode  and  connected  to  the  positive  pole  of  a  series 
of  storage  batteries,  while  the  negative  pole  of  the  series  is  con- 
nected to  the  bearing  in  which  the  shaft  rotates,  thus  allowing 
the  current  to  go  from  the  platinum  plate  through  the  solution 
to  the  crucible,  up  the  shaft  of  the  motor,  and  back  to  the  bat- 
teries. The  power  to  run  the  motor  may  be  conveniently  taken 
from  the  incandescent  light  circuit. 

Experimental  Process:  Copper  and  Nickel.  —  Dissolve  a  weighed  nickel 
five-cent  piece  in  a  mixture  of  10  cm.3  of  sulphuric  acid,  10  cm.3  of  nitric 
acid,  and  30  cm.3  of  water.  Evaporate  on  the  water-bath  or  steam-bath 
to  remove  all  nitric  acid.  Dissolve  the  residue  and  evaporate  again  until 
the  escaping  vapor  contains  no  nitric  acid  as  shown  by  testing  with  litmus 
paper.  Transfer  the  solution  to  a  500-cm.3  flask  and  fill  to  the  mark  with 
water,  mixing  carefully.  Draw  off  50  cm.3  of  the  solution  into  a  150-cm.3 
beaker,  add  three  grams  of  ammonium  sulphate,  dilute  to  about  100  cm.3, 
in  readiness  to  electrolyze  with  the  crucible  used  as  a  rotating  cathode. 
To  make  the  electrolysis  attach  to  the  shaft  of  the  motor  a  weighed  plati- 
num crucible,  to  be  used  as  the  cathode,  in  the  manner  described  on  p.  52 
and  shown  in  Fig.  15.  Adjust  the  levels  so  that  about  two-thirds  of  the 
crucible  shall  be  immersed  in  the  liquid,  exposing  a  cathode  surface  of 
about  30  cm.2.  Introduce  the  anode,  run  the  motor  rapidly  (at  the  rate 
of  about  600  revolutions  per  minute) .  Connect  the  crucible  with  the  nega- 
tive pole  of  the  battery  and  the  anode  with  the  positive  pole.  Electrolyze 
the  solution  with  a  current  of  about  1.5  amp.  to  3  amp.  (corresponding  to 
about  NDioo  =  5  amp.  or  10  amp.).  When  the  deposition  is  nearly  com- 
plete stop  the  rotation;  wash  down  the  sides  of  the  beaker,  the  anode, 
and  the  crucible,  with  a  fine  jet  of  water;  again  start  the  motor,  and  note 
whether  copper  is  deposited  on  the  newly  covered  cathode  surface.  If 
the  deposition  is  complete,  stop  the  motor,  remove  the  crucible  from  the 
liquid,  and  wash  it  with  water,  saving  carefully  the  solution  and  the 
washings.  Then  wash  the  deposit  with  alcohol,  hang  the  crucible  over 
the  end  of  a  nearly  vertical  glass  rod  and  dry  the  deposit  well  above  a 
flame.  Cool  and  weigh.  Calculate  the  percentage  of  copper  in  the  coin. 


54 


QUANTITATIVE  CHEMICAL  ANALYSIS 


Evaporate  the  solution  and  washings  to  a  volume  of  about  20  cm.3, 
neutralize  the  solution  with  ammonium  hydroxide,  add  30  cm.3  of  con- 
centrated ammonium  hydroxide,  filter  if  necessary  and  electrolyze  the  solu- 
tion with  a  current  of  1.5  amp.  to  3  amp.  (corresponding  to  about  NDioo  = 
5  amp.,  or  10  amp.),  depositing  the  nickel  upon  the  crucible  used  as  the 
cathode  and  rotated  at  the  rate  of  about  600  revolutions  per  minute. 
When  the  deposition  is  completed,  stop  the  motor,  remove  the  crucible 
from  the  liquid,  wash  with  water  and  with  alcohol,  dry  carefully  well  above 
a  flame,  cool,  and  weigh.  Remove  the  nickel  from  the  crucible  by  heating, 
for  at  least  fifteen  minutes,  in  nitric  acid  (1  :  1).  State  the  results  in  the 
following  form. 


Weight  of  coin. 

Copper  found. 

Nickel  found. 

Grin. 

Grm. 

Per  cent. 

Grm. 

Per  cent. 

Small  Electrodes  with  Weighable  Cell.  —  Very  small  elec- 
trodes may  be  used  in  electrolytic  analysis,  even  with  high 
currents,  provided  means  be  taken  to  retain  and  gather  that 
portion  of  the  deposit  which  under  such  conditions  is  generally 
detached  from  the  electrode.  The  device  shown  in  Fig.  16  is 
designed  to  accomplish  this  purpose.*  It  consists  of  a  35-cm.3 
cell  C  made  from  a  thistle-tube.  The  anode  a  is  a  small  piece 
of  thin  platinum  foil,  about  a  centimeter  square,  welded  to 
a  platinum  wire  w'  which  is  sealed  into  the  thistle,  as  shown. 
The  cathode  c  consists  of  a  disk  of  gauze  or  foil,  about  a  centi- 
meter in  diameter,  carried  by  a  rod  of  lead  glass  which  is  fused 
into  the  meshes  of  the  gauze  or  into  a  central  hole  in  the  foil, 
an  arrangement  which  gives  great  rigidity  without  the  use  of 
much  platinum.  The  electrical  connection  between  the  disk 
cathode  and  the  shaft  which  carries  the  current  is  made  by  a 
platinum  wire  w.  A  stand  s  of  spring  brass  holds  the  cell 
during  the  process  of  electrolysis,  and  a  similar  lighter  stand 
(best  made  of  sheet  aluminium)  serves  to  support  the  cell  and 
contents  upon  the  balance  pan.  In  order  to  show  details 
clearly,  the  cell  is  represented  in  the  figure  as  resting  lightly 
upon  this  frame.  In  actual  use,  the  cell  is  pushed  downward 
as  far  as  possible,  and  is  held  rigidly. 

In  using  this  apparatus,  the  cathode  cell  is  washed,  dried, 

*  Gooch  and  Burdick:  Am.  Jour.  Sci.,  34,  107  (1912). 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


55 


and  weighed.  The  material  for  analysis,  dissolved  in  about 
15  cm.3  of  liquid,  is  placed  in  the  cell.  To  prevent  possible  loss 
of  liquid  by  spattering,  a  collar  (conical)  cut  from  a  funnel  or 
(concave)  made  from  a  perforated  and  split  watch-glass,  is  set 
upon  the  rim  of  the  thistle.  The  glass  rod  of  the  cathode  is 


6 
FIG.  16. 


fixed  to  the  motor  shaft.  The  electrical  connections  are  made 
as  shown.  The  cathode  is  rotated  at  about  300  revolutions 
per  minute. 

At  the  end  of  the  electrolysis  the  cathode  is  disconnected 
from  the  shaft  and  washed  in  the  cell.  Loose  particles  of  the 
deposit  may  be  allowed  to  settle,  the  supernatant  liquid  de- 
canted, and  after  washing  by  decantation  with  water,  and 
finally  with  alcohol,  the  cell  and  cathode  may  be  dried  in  the 
air-bath,  and  weighed.  The  difficulty  of  washing  thoroughly 
by  decantation  without  loss  of  any  loose  particles,  and  with 
complete  removal  of  the  washing  liquid  so  that  the  drying 
process  need  not  be  prolonged  unduly,  is,  however,  consider- 
able. To  meet  this  difficulty  use  is  made  of  the  little  filter 
tube  shown  in  Fig.  166.  This  is  made  by  fusing  the  end  of  a  lead 
glass  tube,  flared  slightly,  to  a  little  disk  of  platinum  foil  per- 
forated by  a  sharp  point,  the  rough  edges  of  the  perforations 
being  turned  outward.  By  applying  the  suction  of  the  filter 
pump  to  the  open  end  of  the  tube,  and  dipping  the  disk  end 
into  an  emulsion  of  asbestos,  a  felt  of  asbestos  is  deposited 


56  QUANTITATIVE  CHEMICAL  ANALYSIS 

upon  the  perforated  disk.  The  tube  and  felt  may  be  dried, 
and  even  ignited,  without  difficulty. 

This  little  reverse  filter  is  dried  and  weighed  with  the  cell  and 
cathode  ready  for  use.  The  electrolysis  is  made  in  the  manner 
previously  described.  The  liquid  of  the  cell  and  the  washings 
of  the  cathode  and  cell  are  drawn  off  through  the  reverse  filter. 
The  stand,  cell,  cathode,  and  reverse  filter,  shown  in  Fig.  16c, 
are  dried  and  weighed  together. 

In  experimental  determinations  made  with  this  apparatus 
it  has  been  found  that  as  much  as  2  amp.  of  current  may  be 
applied  successfully  to  2  cm.2  of  cathode  surface,  an  amount  of 
current  which  corresponds  to  a  normal  density  of  100  amperes 
(NDioo  =  100  amp.).  Moreover,  fair  results  may  be  obtained 
without  rotation  of  the  cathode,  but  in  this  case  the  deposit  is 
more  bulky,  rather  slow  in  drying,  and  perhaps  more  sensitive 
to  oxidation. 

PROCESSES  DEPENDENT  UPON  DIFFERENCES 
IN  RESPECT  TO   SOLUBILITY 

Processes  of  analysis  frequently  depend  upon  differences  in 
solubility  of  a  substance  in  respect  to  different  liquid  solvents,  and 
upon  differences  between  different  substances  in  respect  to  solubility 
in  the  same  liquid  solvent. 

The  Distribution  of  a  Solute  between  Two  Non-Miscible 
Solvents.  —  A  dissolved  substance  may  sometimes  be  withdrawn 
from  solution  by  agitation  of  the  solution  with  a  more  efficient 
solvent  which  is  not  miscible  with  the  former  solvent.  The  com- 
mon mode  of  identifying  iodine  in  dilute  aqueous  solution  by 
concentrating  it  in  chloroform  or  carbon  disulphide  shaken  up 
with  the  solution  is  a  case  in  point.  In  applying  such  processes 
of  extraction  quantitatively  it  is  well  to  bear  in  mind  that,  in 
accordance  with  the  general  principle  which  governs  the  di- 
vision of  a  solute  between  two  non-miscible  solvents,  the 
efficiency  of  the  operation  increases  rapidly  with  the  number  of 
successive  shakings  with  fresh  portions  of  the  more  efficient 
solvent. 

When  an  amount  of  solute  (zp),  in  the  known  volume  (V) 
of  the  water  solution,  is  shaken  (in  a  separating  funnel)  with  a 
known  volume  (v)  of  the  non-miscible  absorbent,  the  relations 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  57 

of  concentration  of  the  solute  in  water  x\  and  that  of  the  solute 
in  the  absorbent  (XQ  —  x\)  are  given  by  the  expressions  ^  and 

—  respectively.     If  the  ratio  of  the  concentration  of  the 

solute  in  a  cubic  centimeter  of  water  to  its  concentration  in  a 
cubic  centimeter  of  the  absorbent  —  always  a  constant  ratio  — 
be  represented  by  the  distribution  coefficient  k,  then 

Xi        XQ  —  Xi  j  I      kV 

k,   and  Xi  =  sr 


V  v 

The  amount  of  the  solute  (xz)  which  remains  in  the  water  solu- 
tion (V)  after  shaking  it  with  another  similar  portion  of  the 
absorbent  (v)  will  be  given  by  the  equation 

kV 


and  for  n  shakings  with  similar  portions  of  the  absorbent  the 
expression  becomes 

kV 


Xn   =   '. 

Theoretically,  the  absorption  of  the  solute  by  the  more  efficient 
solvent  must  be  always  approximate;  but,  practically,  it  is  often 
possible  to  arrange  the  conditions  so  that  the  amount  of  solute 
which  escapes  absorption  will  be  insignificant. 

The  extraction  of  iodine  from  a  water  solution  by  shaking 
with  carbon  disulphide  *  (or  chloroform)  and  the  separation  of 
iron  and  molybdenum  from  nickel,  vanadium,  and  chromium 
by  the  ether  extraction  of  the  aqueous  solution  of  the  chlorides  f 
illustrate  the  application  of  this  principle. 

Precipitates  and  Conditions  Affecting  Precipitation 

The  exactness  of  processes  dependent  upon  differences  between 
substances  in  respect  to  solubility  is  determined  by  the  attainment 
of  conditions  conducive  to  the  highest  degree  of  insolubility,  purity, 
and  definiteness  of  precipitates.  With  these  conditions  assured 
to  a  reasonable  degree,  the  practical  separation  of  the  precipitate 
from  the  medium  in  which  it  is  formed  also  requires  consideration. 

*  See  p.  185. 

t  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed.,  Vol.  II,  p.  313. 


58  QUANTITATIVE  CHEMICAL  ANALYSIS 

Conditions  Affecting  the  Insolubility  of  Precipitates.  — As 

has  been  previously  shown  *  the  velocity  (if  measurable)  with 
which  a  non-reversible  reaction  is  completed  and  the  degree  to 
which  a  reversible  reaction  approaches  completion  depend  upon 
the  adjustment  of  the  concentrations  "of  the  reagents  which 
enter  into  action.  In  precipitation  processes  it  is  important, 
therefore,  to  give  due  attention  to  the  concentrations  of  the  re- 
agents, but  other  conditions  affecting  the  insolubility  of  the 
precipitates  must  likewise  be  taken  into  account,  such  as  the 
amount  and  nature  of  the  medium  from  which  the  precipitate 
separates  and  the  state  of  subdivision  in  which  it  exists  in  the 
medium. 

The  Nature  of  the  Medium.  —  Media  in  which  precipitation 
may  take  place  may  differ  in  respect  to  the  solvent-effects 
which  they  are  able  to  exert  upon  the  precipitate.  The  in- 
solubility of  a  precipitate  may  sometimes  be  increased  ad- 
vantageously by  adding  to  the  medium  in  which  precipitation 
is  effected  some  miscible  liquid  in  which  the  precipitate  is  more 
insoluble  than  it  is  in  the  original  medium.  For  example,  the 
precipitation  of  substances  which  are  more  insoluble  in  alcohol 
than  in  water  may  be  made  more  complete  by  adding  alcohol 
to  the  water  solution.  For  this  reason,  potassium  fluosilicate 
and  calcium  sulphate  being  more  insoluble  in  the  alcoholic  mix- 
ture than  in  pure  water,  it  is  customary  to  mix  alcohol  with  the 
water  solution  of  potassium  chloride  from  which  potassium 
fluosilicate  is  to  be  precipitated  by  fluosilicic  acid  and  with 
solutions  of  calcium  chloride  from  which  calcium  sulphate  is  to 
be  precipitated  by  sulphuric  acid.  In  such  cases  the  volume  of 
the  alcoholic  liquid  must  not  be  made  too  great,  because  the  pre- 
cipitates are  not  absolutely  insoluble  even  in  that  medium  and 
the  concentration  of  the  alcohol  should  not  be  great  enough  to 
induce  the  precipitation  of  substances  which  should  remain  in 
solution. 

The  State  of  Subdivision.  —  A  very  finely  divided  crystalline 
substance  tends  to  exhibit  greater  solubility  than  does  the 
same  substance  when  it  exists  in  coarser  particles  under  other- 

*  See  p.  5. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  59 

wise  similar  conditions.  For  example,  it  is  known  that  calcium 
sulphate  has  a  definite  solubility  of  2.085  grm.  in  1  liter  of 
water  at  25  degrees  (approximately,  1  : 480),  provided  the 
saturated  solution  is  in  contact  with  particles  of  diameter 
greater  than  0.002  mm.,  while  the  same  substance  in  particles 
averaging  0.0003  mm.  (and  not  less  than  0.0002  mm.  in  diameter) 
goes  into  solution  to  the  extent  of  2.476  grm.  to  1  liter  (approxi- 
mately 1  : 400)  at  25  degrees.  Similarly,  barium  sulphate 
dissolves  in  the  proportion  of  0.00229  grm.  to  1  liter  of  water 
(1  : 400,000)  when  the  diameter  of  the  particles  approximates 
0.018  mm.  and  in  the  proportion  of  0.00415  grm.  to  1  liter 
(1  : 250,000)  when  the  diameter  of  the  particles  is  less  than 
0.00001  mm.  If  larger  particles  of  a  substance  are  placed  in  the 
solution  saturated  in  contact  with  finer  particles,  the  equi- 
librium between  the  solid  and  the  solution  is  disturbed  and 
the  substance  in  solution  is  deposited  upon  the  larger  grains, 
while  the  smaller  grains  dissolve  to  keep  up  the  saturation  cor- 
responding to  the  solubility  of  the  finer  particles  until  they 
finally  disappear.  Thus,  very  minutely  crystalline  precipitates 
tend  to  become  more  coarsely  crystalline  upon  long  digestion  in 
the  solutions  from  which  they  separate,  the  solubility  diminish- 
ing at  the  same  time.  A  substance  formed  in  solution  in  a 
state  of  subdivision  corresponding  to  a  certain  degree  of  solu- 
bility may,  however,  remain  dissolved  to  a  degree  which  would 
correspond  to  supersaturation  with  respect  to  the  same  sub- 
stance in  coarser  and  more  insoluble  state  of  subdivision  until  a 
change  of  condition,  such  as  change  of  temperature  or  surface 
evaporation  on  standing,  causes  the  incipient  formation  of  larger 
particles,  which  once  started,  will  proceed  until  an  equilibrium  is 
established  between  the  substance  in  solution  and  the  substance 
in  larger  particles.  Such  is  probably  the  state  of  affairs  when 
precipitates,  like  calcium  sulphate,  or  even  the  ordinarily  very 
insoluble  barium  sulphate,  are  deposited  tardily  from  dilute 
solutions.  The  variation  of  solubility  with  the  degree  of  sub- 
division of  crystalline  substances,  with  the  consequent  growth 
of  larger  particles  at  the  expense  of  smaller  ones,  appears  to  be 
a  result  of  struggle  toward  equilibrium  between  the  forces  of 


60  QUANTITATIVE  CHEMICAL  ANALYSIS 

cohesion  and  tension  acting  at  the  surface  of  contact  between  the 
solid  and  the  liquid. 

Colloidal  Suspensions.  —  In  colloidal  suspension,  certain 
substances,  like  the  acidic  hydroxides  of  silicon,  titanium,  and 
tungsten,  or  the  basic  hydroxides  of  iron,  aluminium,  and 
chromium,  or  the  sulphides  of  arsenic  and  tin,  appear  to  consist 
of  particles  in  a  much  finer  state  of  subdivision  than  that  of  the 
finest  crystalline  precipitates.  While  the  diameters  of  the  finest 
crystalline  particles  are  of  the  order  of  thousandths  of  a  milli- 
meter, those  of  the  colloidal  suspensions  —  invisible  with  the 
ordinary  arrangement  of  -the  microscope  and  only  visible  when, 
like  motes  in  a  sunbeam,  they  are  observed  by  the  microscope 
placed  perpendicular  to  a  strong  beam  of  light  illuminating  them 
—  appear  to  possess  diameters  of  the  order  of  millionths  of  a 
millimeter.  It  seems  probable  that,  like  variations  in  the  solu- 
bility of  crystalline  salts  with  degree  of  subdivision,  the  peculiar 
properties  of  colloidal  substances  have  relation  to  fineness  of 
subdivision,  or  to  extent  of  contact  surface  in  comparison  with 
mass. 

In  presence  of  electrolytes,  some  colloidal  substances  may  ap- 
pear as  coagulated  hydrogels,  while  in  water  sufficiently  devoid 
of  such  coagulating  agents  they  may  take  the  form  of  hydrosols 
in  suspension  or  pseudosolution.  The  proportion  of  the  elec- 
trolyte required  to  produce  coagulation  may  be  exceedingly 
small  but  it  must  exceed  a  certain  limit,  and  free  acids  and  salts 
of  polyvalent  elements  (easily  susceptible  to  hydrolytic  action) 
are  most  effective,  the  coagulating  power  of  these  electrolytes 
at  similar  concentration  increasing  very  markedly  with  the  va- 
lence of  the  cation.  In  the  process  of  coagulation,  however,  the 
precipitate  takes  up  contaminating  material  derived  from  the 
coagulant,  and  in  analytical  operations  it  is  therefore  usual  to 
employ  volatile  acids  or  ammonium  salts  of  volatile  acids  as 
coagulants,  in  order  that  the  contamination  of  the  colloidal 
precipitates  may  be  subsequently  removed  by  volatilization. 
When,  however,  fixed  contaminating  material  has  been  intro- 
duced into  a  colloidal  precipitate  it  may  be  replaced,  more  or 
less  completely,  by  the  action  of  another  salt,  and  may  often  be 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  61 

removed  from  the  colloidal  precipitate  by  digestion  and  wash- 
ing with  ammonium  salts. 

In  analytical  practice  it  is  not  uncommon  to  note  that  a  col- 
loidal precipitate  which  is  insoluble  in  the  presence  of  suitable 
coagulating  agents  will  when  washed  with  pure  water  pass 
through  the  filter  in  clear  solution  or  in  turbid  pseudosolution; 
and  that  from  such  pseudosolutions  or  suspensions  the  hydrogel 
is  again  precipitated  by  the  action  of  the  coagulating  agents  in 
the  filtrate. 

It  is  to  be  noted  also  that  the  colloidal  nature  of  a  precipi- 
tate may  often  be  modified  by  evaporating  the  suspending 
liquid  and  drying  the  residue  (e.g.,  silica),  and  in  some 
cases  by  simply  digesting  the  precipitate  (e.g.,  alumina) 
in  the  heated  liquid. 

Purity  of  Precipitates.  —  All  precipitates  tend,  more  or  less 
according  to  their  nature,  to  carry  down  with  them  other  sub- 
stances from  the  solution.  This  tendency  is  least  in  the  case 
of  coarsely  granular  substances,  greater  on  the  part  of  mi- 
nutely granular  substances,  and  greatest  in  the  case  of  colloidal 
precipitates.  Whenever  it  is  a  possibility,  processes  should  be 
employed  which  produce  precipitates  of  crystalline  habit  and 
care  should  be  taken  to  secure  a  proper  coarseness  of  grain  by  the 
choice  of  suitable  conditions  as  regards  the  medium,  the  con- 
centrations of  reagents,  the  temperature,  and,  if  necessary, 
suitable  digestion.  When  colloidal  precipitates  are  necessarily 
put  to  use  in  analytical  processes  they  should  be  modified  as 
much  as  may  be  by  digestion  at  a  suitable  temperature  or  by 
drying  and,  if  they  are  to  be  weighed,  they  should  be  washed, 
if  possible,  with  suitable  reagents  which  may  later  be  volatilized, 
such  as  ammonium  salts  of  volatile  acids  to  replace  adsorbed 
material.  In  some  cases  they  must  be  redissolved  and  further 
purified  by  reprecipitating  and  finally  washing  with  suitable 
reagents  which,  though  adsorbed,  may  be  later  volatilized. 
In  all  processes  in  which  the  addition  of  the  precipitant  brings 
about  immediate  precipitation  the  precipitating  reagent  should 
be  added  slowly  to  diminish  the  chances  of  mechanical  inclu- 
sion of  foreign  matter  in  the  precipitate. 


62  QUANTITATIVE  CHEMICAL  ANALYSIS 

Definiteness  in  Composition  of  Precipitates.  —  Wholly 
apart  from  the  effects  of  contamination  by  the  adsorption  of 
impurities  taken  from  the  solution,  precipitates  sometimes 
exhibit  variations  from  an  ideal  chemical  composition  in  con- 
sequence of  secondary  reactions  with  other  substances  than 
those  concerned  in  the  reaction  upon  which  the  analytical 
process  is  based  or  because  changes  in  the  concentrations  of 
factors  in  the  reaction  make  changes  in  the  composition  of  prod- 
ucts. For  example,  care  must  be  taken  in  precipitating 
ammonium  magnesium  phosphate  or  ammonium  manganese 
phosphate  to  control  the  excess  of  ammonium  salt  also  present 
in  order  that  the  precipitate  may  have  the  ideal  composition, 
and  yield  on  ignition  the  ideal  pyrophosphate.* 

The  Practical  Separation  of  Precipitates.  —  Filtering  and 
Washing  of  Precipitates.  In  separating  precipitates  by  decan- 
tation  of  the  liquid  in  which  they  are  formed,  or  by  filtration 
of  the  liquid  through  paper,  asbestos,  platinum  sponge,  or  other 
suitable  material,  it  is  necessary  to  overcome  the  viscosity  of 
the  liquid  and  the  attraction  between  the  insoluble  solid  and 
the  liquid.  In  this  process,  the  residual  or  adherent  solution 
must  be  removed  from  the  precipitate  by  washing  with  an- 
other miscible  liquid  inactive  toward  the  precipitate  and  sep- 
arable from  it  by  volatilization  at  a  temperature  which  does 
'not  endanger  its  stability  and  definiteness. 

Upon  the  presumption  that  the  residual  solution  and  the 
washing  liquid  mix  homogeneously  in  every  filtration  or  de- 
cantation,  the  washing  is  effective  in  proportion  to  the  relation 
between  the  total  volume  (V)  of  the  liquid  before  decanting  or 
filtering  and  the  residual  volume  (v)  left  after  decanting  or 
filtering.  After  a  single  washing,  the  material  remaining  is 

then  a  definite  fraction  y  of  the  material  contained  in  the 

residual  solution  before  the  washing;  and  the  effect  of  (n)  suc- 
cessive similar  washings  will  be  given  by  the  corresponding 


power  of  the  fraction  —  viz.,  f^J   .     Thus,  if  the  ratio  of  the 
residual  liquid  to  the  total  liquid  is  uniformly  J,  a  single  wash- 

*  See  p.  81. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  63 

ing  reduces  the  material  in  solution  to  |,  two  washings  to  J,  and 

(1\4       1 
-J   =  ^  of  the  original  amount.     When  the 

total  volume  is  ten  times  as  large  as  the  volume  of  the  residual 
liquid,  the  same  number  of  washings  reduces  the  material  to 

(1  \4          1 
To)  ==  loooo  of  the  original  amount- 

It  is  advantageous  to  make  the  volume  of  the  liquid  de- 
canted or  drained  from  a  filter  large  in  comparison  with  the 
volume  of  the  residual  liquid;  and  similarly,  it  is  best  to  drain 
off  in  the  decantation  or  filtration  as  much  liquid  as  possible 
before  adding  a  new  portion  of  the  washing  liquid.  In  practice, 
however,  washing  is  not  usually  effected  with  the  rapidity  and 
thoroughness  which  might  be  indicated  by  a  consideration  of 
the  amount  and  number  of  the  portions  of  washing  liquid 
used;  for  when  porous  solids  include  solutions  homogeneous 
mixing  of  the  solutions  with  an  external  washing  liquid  takes 
place  only  by  the  slow  process  of  diffusion,  and  many  precipi- 
tates show  a  very  marked  tendency  to  hold  by  adsorption 
matter  which  is  dissolved  in  liquids  in  contact  with  them. 
The  efficiency  of  the  washing  process  is,  therefore,  best  de- 
termined by  chemical  tests  applied  to  the  filtrate.  In  practical 
work  the  conditions  of  precipitation  must  be  regulated  in- 
telligently for  each  individual  case  to  insure  insolubility  or  the 
nearest  approach  to  insolubility,  the  highest  degree  of  freedom 
from  adsorbed  or  mechanically  included  foreign  matter,  and 
favorable  conditions  for  washing. 

Filtration  by  Paper.  —  The  paper  filter  used  in  the  separation 
of  precipitates  intended  for  weighing  should  be  "ashless" 
that  is,  of  the  (purchaseable)  sort  which  has  been  washed  with 
hydrochloric  acid  and  with  hydrofluoric  acid,  and  which  leaves 
on  ignition  a  very  small  known  weight  or  an  inappreciable 
weight.  The  disk  of  paper  is  folded  to  a  quadrant  and  opened 
to  a  cone  of  an  angle  a  trifle  larger  than  that  of  the  funnel 
which  it  is  to  fit.  The  funnel  should  have  an  angle  of  60  degrees 
and  a  long  thin  stem.  The  paper,  held  within  the  funnel  by  a 
finger,  is  filled  with  water,  and  patted  into  close  contact  with  a 


64  QUANTITATIVE  CHEMICAL  ANALYSIS 

glass,  with  special  care  to  have  the  fitting  perfect  on  the  upper 
rim.  Whenever  it  is  possible,  it  is  very  desirable  to  settle  the 
precipitate  carefully  and  to  decant  upon  the  filter  the  clear  su- 
pernatant liquid  to  the  lowest  possible  limit;  and  it  is  frequently 
desirable  to  repeat  the  washing  by  dilution  and  decantation. 
The  precipitate,  transferred  from  the  container  to  the  filter  (with 
care  to  leave  margin  of  paper  half  a  centimeter  broad  above  the 
level  of  the  liquid)  by  means  of  a  fine  jet  of  water  directed  from 
a  wash-bottle,  and  with  the  aid  of  a  rubber-tipped  glass  rod 
(the  "policeman")  to  remove  any  of  the  precipitate  which  may 
adhere  to  the  walls  of  the  container,  is  washed  down  into  the 
point  of  the  filter  with  a  rotary  motion  of  the  water  jet.  The 
washing  should  be  continued  until  suitable  tests  show  that  the 
precipitate  is  not  yielding  foreign  material  to  a  fresh  washing. 
The  filter,  with  the  precipitate,  is  usually  dried,  at  least  partially, 
and  folded  compactly,  with  the  precipitate  inside,  in  readiness 
to  be  ignited  in  a  weighed  crucible.  Sometimes,  after  drying, 
the  precipitate  and  paper  are  separated  so  far  as  is 
possible  and  ignited  apart. 

If  the  precipitate  is  not  colloidal,  and  not  sus- 
ceptible to  the   compacting  action  of  pressure,   a 
vacuum  flask  or  bell-jar,  evacuated  by  a  filter  pump 
FIG    17      (Fig-   17),  may  be  advantageously  used  to  receive 
the  filtrate;    but  in  this  case  it    is    necessary    to 
strengthen  the  point  of  the  filter  by  placing  between  it  and  the 
funnel  a  cone  of  perforated  .platinum  or  of  parchment  paper, 
so  that  it  will  withstand  the  atmospheric  pressure  from  above. 
Filtration  by  Asbestos.  —  In  many  processes,  such  as  those  in 
which  the  precipitate  is  affected  in  ignition  by  the  presence  of 
carbonaceous  matter  or  in  which  the  precipitate  must  be  dried 
rather  than  ignited,  asbestos  used  in  the  filtering  crucible* 
may  be  advantageously  substituted  for  the  paper  filter.     White, 
silky,  flexible,  anhydrous,  asbestos  fiber  f  is  disintegrated,  best 
by  scraping  longitudinally  to  a  fine  soft  down,  and  prepared 

*  Gooch:   Proc.  Amer.  Acad.,  13,  342  (1878);  Am.  Chem.   Jour.,  1, 
317. 

f  The  serpentine  variety  is  hydrous  and  useless  for  this  purpose. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  65 

for  use  by  boiling  with  hydrochloric  acid  and  washing  with 
water.  In  the  first  stage  of  the  washing  it  is  best  to  stir  the 
asbestos  into  a  large  volume  of  water  and  to  allow  the  coarser 
pieces,  unsuitable  for  making  the  pulp,  to  settle  out.  Then 
the  main  portion  of  the  suspended  asbestos  is  allowed  to  settle, 
the  supernatant  very  finest  particles  are  poured  off  into  an- 
other receptacle,  and  the  washing  may  be  completed,  with  the 
aid  of  the  filter  pump,  on  a  perforated  cone  or  in  a  draining 
funnel  lined  with  a  paper  disk.  If  the  fine  material  is  left  with 
the  rest,  the  filter  made  from  the  mixture  is  likely  to  be  too 
close  and  compact  for  the  rapid  filtration  which  is  desirable 
and  perfectly  feasible  in  ordinary  work;  but,  if  preserved  by 
itself  it  may  be  used  to  supple- 
ment the  more  porous  material 
when  a  very  compact  filter  is 
required. 

The  filter  is  made  in  a  crucible 
(platinum  or  porcelain)  the  bot- 
tom of  which  is  perforated  with 
fine  holes,  the  more  the  better. 
The  crucible  is  held  by  a  rubber 
setting  made  by  stretching  soft 
rubber  tubing  over  a  funnel,  or  thistle  tube,  as  shown  in 
Fig.  18a,  and  the  funnel  is  fitted  tightly  in  one  of  the 
perforations  of  the  stopper  which  closes  the  vacuum  flask.  A 
little  of  the  asbestos  is  suspended  in  water,  poured  into  the 
crucible,  while  the  vacuum  pump  is  in  action,  takes  the  form 
of  a  compact  felt,  which  is  easily  washed,  dried,  ignited,  and 
weighed  with  the  crucible.  When  a  precipitate  is  to  be  filtered 
off  and  weighed  upon  a  felt  thus  prepared,  the  crucible,  with  the 
felt  in  place,  is  placed  in  the  rubber  setting,  the  asbestos  is 
moistened  with  a  drop  or  two  of  water,  the  pump  is  started, 
and,  after  the  crucible  has  settled  into  place  (see  Fig.  186), 
the  filtration  and  washing  are  carried  on  in  the  usual 
manner. 

The  Filter  of  Spongy  Platinum.  —  Instead  of  asbestos  spongy 
platinum  may  sometimes  be  used  with  advantage,  in  the  per- 


66  QUANTITATIVE  CHEMICAL  ANALYSIS 

f orated  crucible,  as  the  filtering  medium.*  The  felt  may  be 
made  by  placing  the  perforated  platinum  crucible  upon  ab- 
sorbent paper,  pouring  into  the  crucible  a  thick  emulsion  of  am- 
monium chloroplatinate,  and  igniting  while  the  layer  of  the 
resulting  spongy  platinum  is  shaped  by  pressure  with  a  rounded 
glass  rod;  or  platinum  crucibles  with  porous  bottoms  (Neu- 
bauer)  may  now  be  procured  from  the  manufacturers. 

The  Anthracene  Filter.  —  In  some  processes  of  analysis  the  necessity 
arises  for  the  separation  of  a  precipitate  from  the  filter  upon  which  it  has 
been  collected  or  of  mixed  precipitates  from  one  another.  When  a  precipi- 
tate yields  easily  to  a  solvent  which  does  not  affect  paper  the  use  of  the 
paper  filter  to  collect  the  precipitate  to  be  subsequently  dissolved  is  feasi- 
ble. When  the  solvent  for  the  precipitate  attacks  paper  also,  as  is  the  case 
with  the  stronger  acids  and  alkali  hydroxides  in  concentrated  solution,  so 
that  the  separation  of  the  solution  from  the  disintegrated  paper  makes 
trouble,  the  asbestos  filter  may  be  a  serviceable  substitute  for  the  paperfilter. 
If,  however,  the  separation  of  a  refractory  precipitate  from  the  filter  upon 
which  it  has  been  collected  is  required,  the  problem  may  sometimes  be 
solved  satisfactorily  by  the  use  of  the  anthracene  filter  f  insoluble  in  water 
and  aqueous  solutions  of  ordinary  salts,  alkalies,  and  acids  (excepting 
strong  sulphuric,  strong  nitric,  and  glacial  acetic  acid),  but  easily  soluble 
in  naphtha,  benzene,  carbon  disulphide,  and  in  boiling  alcohol. 

The  mode  of  preparing  and  using  the  filter  is  simple.  Anthracene  is 
slightly  moistened  with  alcohol  so  that  water  will  wet  it.  An  emulsion  in 
water  is  then  made  and  this  is  used  in  the  same  way  as  is  the  emulsion  of 
asbestos  employed  in  making  asbestos  felts.  That  is  to  say,  enough  of 
the  emulsion  to  form  a  layer  of  the  proper  thickness  is  poured  into  a  per- 
forated crucible  which  is  held  tightly  in  the  rubber  packing  of  a  funnel 
fitted  in  the  usual  manner  to  a  vacuum-flask.  If  the  felt  happens  to  be 
too  coarse  for  a  special  use,  it  may  be  made  as  close  as  need  be  by  coating 
the  layer  first  deposited  with  a  finer  emulsion,  made  by  dissolving  anthracene 
in  hot  alcohol  and  precipitating  with  water.  As  a  protection  against  disin- 
tegration of  the  felt  of  anthracene  in  the  subsequent  processes  of  filtration 
and  washing  a  perforated  disk  of  platinum  foil  of  diameter  a  little  less  than 
that  of  the  crucible  may  be  placed  upon  the  felt.  After  washing  with 
water  the  filter  is  ready  for  use. 

To  separate  a  precipitate  from  the  anthracene  filter  it  is  only  necessary 
to  act  with  a  solvent  for  anthracene.  It  is  usually  convenient  to  stand 
the  crucible  containing  precipitate  and  filter  in  a  small  beaker,  add  enough 
of  the  solvent  for  anthracene  and  gently  warm  until  the  anthracene  dis- 

*  Munroe:  J.  Anal.  Chem.,  2,  241;   Chem.  News,  58,  101. 
t  Gooch:  Proc.  Amer.  Acad.,  20,  390  (1885). 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  67 

solves.  On  the  addition  of  a  water  solution  of  the  reagent  to  act  upon  the 
precipitate  the  solution  of  anthracene  floats  to  the  surface  of  the  water 
solution.  If  the  precipitate  dissolves  entirely,  the  solution  of  anthracene 
may  be  separated  from  the  water  solution  by  simply  pouring  the  liquids 
upon  a  filter  previously  moistened  with  water,  when  the  solution  in  water 
will  run  through,  and  the  anthracene  and  its  solvent  will  remain  and  may 
be  washed  indefinitely  with  water.  In  case  a  constituent  of  a  mixture  of 
precipitates  dissolves,  and  it  is  desired  to  collect  another  which  is  insoluble, 
the  anthracene  and  its  solvent  may  be  made  to  pass  the  filter,  after  the 
water  has  run  through,  by  adding  a  little  alcohol  to  overcome  the  repulsion 
between  the  solution  and  the  water  which  fills  the  pores  of  the  filter.  The 
precipitate  which  stays  behind  may  then  be  washed  first  with  a  solvent  of 
anthracene,  and  then,  if  necessary,  with  alcohol  followed  by  water.  If 
the  vacuum  filter  is  used  (with  either  paper  or  asbestos)  both  liquids, 
the  aqueous  solution  and  the  immiscible  solution  of  anthracene,  will  leave 
the  precipitate  and  traverse  the  filter  together. 

In  general,  benzol  is  to  be  preferred  as  the  solvent  for  anthracene,  but 
some  advantage  may  be  gained  in  special  cases  by  a  proper  choice  of 
solvents.  Thus,  in  removing  intermixed  sulphur  from  precipitated  sul- 
phides collected  upon  an  anthracene  filter  both  the  anthracene  and  the 
sulphur  may  be  dissolved  in  carbon  disulphide  in  a  single  operation. 

Representative  Precipitation  Processes.  —  Certain  pro- 
cedures involving  the  formation  of  chlorides,  sulphates,  car- 
bonates, oxalates,  phosphates,  hydroxides,  and  sulphides  for  the 
purpose  of  determining  either  acidic  ions  or  basic  ions  of  re- 
action may  be  taken  as  fairly  typical  in  respect  to  the  phenom- 
ena and  the  treatment  incidental  to  precipitation  processes. 

The  Precipitation  of  Chlorides 

The  Determination  of  the  Chloride  Ion.  —  When  silver 
nitrate  is  added  to  a  dilute  solution  containing  hydrochloric 
acid  or  a  chloride,  silver  chloride  is  formed  and  nearly  all  of 
it  is  precipitated. 

AgNOa  +  HCl   <=»AgCl  +  HN03. 

AgNO3  +  NaCl  <=±  AgCl  +  NaNO3. 

The  freshly  precipitated,  finely  divided,  and  somewhat  col- 
loidal silver  chloride  becomes  more  compact  and  somewhat  less 
soluble  upon  heating  or  upon  long  standing  in  the  solution,  and 
the  presence  of  even  a  slight  excess  of  silver  nitrate  reduces  the 


68  QUANTITATIVE  CHEMICAL  ANALYSIS 

solubility  to  an  amount  practically  negligible  in  the  volumes 
used  in  ordinary  analytical  processes.  Sodium  nitrate  and,  to 
a  less  extent,  nitric  acid,  though  useful  in  promoting  the  coagu- 
lation of  the  precipitate,  may,  if  present  in  considerable  amounts, 
appreciably  increase  the  solubility  of  the  precipitate.  In  a 
liter  of  water  containing  8.5  grm.  (TV  gram-molecule)  of  sodium 
nitrate  as  much  as  0.0030  grm.  of  silver  chloride  may  be  dis- 
solved. Other  nitrates,  notably  mercuric  nitrate,  have  similar 
effects.  Silver  chloride  possesses  in  considerable  degree  the 
tendency  to  occlude  substances  present  in  solution  during  its 
formation,  and,  as  is  natural,  this  tendency  increases  with  in- 
crease in  the  concentrations  of  the  dissolved  substances.  From 
the  freshly  formed,  curdy,  and  flocculent  precipitate  the  occluded 
material  may  be  easily  removed  by  washing,  but  it  is  much 
more  difficult  to  thus  purify  the  precipitate  which  has  been 
compacted  by  long  standing  in  presence  of  the  mother  liquor. 
In  precipitating  silver  chloride  in  the  analytical  process  the 
precipitate  is  subject  to  contamination  with  silver  nitrate  when 
this  reagent  is  in  excess,  but  if  the  precipitations  are  made  in 
solutions  containing  no  more  than  £  gram-molecule  of  the 
soluble  chloride,  by  the  addition  of  silver  nitrate  of  similar 
concentration,  the  material  left  after  washing  proves  to  be  in- 
significant when  the  washing  quickly  follows  the  precipitation. 
If,  however,  the  precipitate  is  to  be  made  compact  by  standing 
for  a  long  time  in  contact  with  the  solution  both  the  solution 
and  the  reagent  should  be  of  concentrations  not  exceeding 
TV  gram-molecule  in  the  liter.  For  example,  in  the  reaction  be- 
tween sodium  chloride  and  silver  nitrate  the  reagents  should 
react  in  solutions  containing  not  more  than  0.58  and  1.7  grm.  of 
these  substances,  respectively,  in  every  100  cm.3  of  solution,  if 
the  precipitate  is  to  stand  long  in  contact  with  the  solution. 

The  precipitate  is  best  collected  on  asbestos  in  the  per- 
forated crucible,  carefully  washed,  dried  at  130  to  140  degrees 
and  weighed. 

It  is  also  possible  to  collect  the  precipitate  upon  ashless 
paper.  In  this  case  the  paper  and  precipitate  are  dried,  and  the 
precipitate  is  separated  as  completely  as  possible  from  the 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


69 


paper  and  reserved.  The  paper  is  carefully  rolled  up  and 
burned  in  a  spiral  of  platinum  wire  over  a  weighed  crucible 
into  which  the  ash  is  allowed  to  fall.  The  ash  is  treated  with 
a  drop  or  two  of  nitric  acid  (to  dissolve  reduced  silver)  and  then 
with  a  drop  or  two  of  hydrochloric  acid  (to  form  silver  chloride 
again),  the  main  portion  of  the  precipitate,  previously  sepa- 
rated, is  added  and  the  entire  amount  of  silver  chloride  is 
heated  to  incipient  fusion,  and  weighed. 

Experimental  Process:  the  Chloride  Ion  in  Barium  Chloride.  —  Weigh  out 
carefully  about  0.5  grm.  of  barium  chloride,  BaCl2  •  2  H^O.  Dissolve  the 
weighed  material  in  about  300  cm.3  of  water.  Add  a  few  drops  of  nitric 
acid  and  then  a  solution  of  silver  nitrate,  gradually  and  with  constant 
stirring  to  curd  the  precipitate,  until  no  further  precipitation  takes  place. 
Heat  the  liquid  to  boiling  and  set  the  whole  aside  in  the  dark  to  cool. 
When  the  precipitate  has  settled  completely  filter  it  off  on  asbestos  in  a 
perforated  crucible  previously  weighed,  wash  thoroughly  with  water  con- 
taining a  few  drops  of  nitric  acid  to  the  liter  (to  keep  the  precipitate  from 
passing  into  the  colloidal  condition)  and  then  with  a  little  pure  water.  Dry 
to  constant  weight,  first  at  100  degrees  and  then  in  an  air-bath  heated  to 
at  least  130  degrees  or  over  a  low  flame  at  a  temperature  well  below  that 
of  the  fusing  point  of  silver  chloride  (450  degrees).  Calculate  the  weight  of 
chlorine  in  the  silver  chloride  weighed  and  compare  it  with  the  weight  which 
should  be  contained  in  the  amount  of  barium  chloride  taken.  Tabulate 
the  results  of  analysis. 


BaCl2  •  2  H2O 
taken. 

AgCl 
formed. 

Cl  in  AgCl 
formed. 

Cl  present  in 
BaCl2  •  2  H2O  taken. 

Error  in  terms  of  Cl. 

Grm. 

Grm. 

Grm. 

Grm. 

Per  cent. 

Grm. 

Per  cent. 

Bromide  Ion  and  Iodide  Ion.  —  The  bromide  ion  and  the 
iodide  ion  of  soluble  salts  may  be  determined  similarly,  by  pre- 
cipitating with  silver  nitrate  and  weighing  the  insoluble  silver 
salt  collected  on  asbestos. 

The  Determination  of  the  Silver  Ion.  —  The  gravimetric 
determination  of  silver  may  be  accomplished  by  taking  ad- 
vantage of  the  reaction  between  a  soluble  silver  salt  and  a 
soluble  chloride,  hydrochloric  acid,  which  is  easily  removed  by 


70  QUANTITATIVE  CHEMICAL  ANALYSIS 

washing  and  is  volatile  in  the  process  of  drying,  being  the  best 
.  precipitant. 

HC1  +  AgN03  <=±  AgCl  +  HN03. 

The  solubility  of  the  silver  chloride  is  less  at  low  temperatures 
and  is  diminished  by  the  presence  of  hydrochloric  acid  in 
moderate  excess,  although  considerably  increased  by  the 
action  of  very  large  amounts  of  that  acid,  or  of  its  soluble 
salts.  While  100  cm.3  of  pure  water  will  dissolve  0.000154  grm. 
of  silver  chloride  at  20  degrees  (and  0.00217  at  100  degrees), 
the  solubility  of  the  silver  chloride  amounts  at  21  degrees  to 
0.00002  grm.,  0.00003  grm.,  and  0.00555  grm.  in  100  cm.3  of 
hydrochloric  acid  of  1,  5,  and  10  per  cent,  respectively. 

Experimental  Process:  Determination  of  Silver  as  Silver  Chloride.  — 
Make  a  nearly  N/10  solution  of  silver  nitrate  by  dissolving  17  grm.  of 
silver  nitrate  in  a  liter  of  water.  Of  this  solution,  draw  from  a  burette  an 
exactly  measured  portion  (50  cm.3).  Add  a  few  drops  of  nitric  acid  (for 
its  favorable  influence  upon  the  condition  of  the  precipitate),  dilute  the  so- 
lution to  a  volume  of  about  200  cm.3  with  boiling  water  (so  that  the  tem- 
perature may  be  at  least  70  degrees),  add  hydrochloric  acid  slowly,  with 
vigorous  stirring  (to  curd  the  precipitate)  until  the  precipitation  appears  to 
be  complete.  Set  the  mixture  aside  to  cool  and,  after  the  precipitate  has 
settled  leaving  the  supernatant  liquid  clear,  test  with  a  drop  or  two  of  hy- 
drochloric acid  to  make  sure  that  the  reaction  has  been  completed.  Filter 
(preferably)  on  asbestos  in  the  perforated  crucible  (previously  ignited  and 
weighed)  and  wash  with  cold  water.  Dry  the  precipitate  upon  the  filter 
to  constant  weight  at  140  degrees,  or  (cautiously)  over  a  very  low  flame  at 
a  temperature  below  the  fusing  point  of  silver  chloride  (450  degrees). 

Or  the  precipitate  may  be  collected  on  papier  and  treated  subsequently 
as  described  in  the  process  for  the  determination  of  the  chlorine  ion  by 
precipitation  as  silver  chloride  (p.  69). 

Calculate  the  silver  content  of  the  solution  per  cm.3. 

Determination  of  Mercury  as  Mercurous  Chloride.  —  The  de- 
termination of  mercury  in  the  mercurous  form  by  precipitation 
as  mercurous  chloride,  Hg2Cl2,  may  be  made  by  treating  the 
solution  of  mercurous  nitrate  with  hydrochloric  acid,  filtering 
and  washing  on  asbestos  in  the  perforated  crucible,  after  stand- 
ing for  twelve  hours,  and  drying  at  105  degrees: 

.  Hg2(NO3)2  +  2  HC1  +±  Hg2Cl2  +  2  HN03. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  71 

The  same  process  may  be  applied  to  the  analysis  of  mercuric 
nitrate  if  phosphorous  acid  be  also  added  to  effect  a  preliminary 
reduction : 

2  Hg(N03)2  +  H3P03  +  H20  =  Hg2(N03)2  +  H3P04  +  2  HN03. 

Determination  of  Barium  as  Chloride.  —  Water  which  dis- 
solves barium  chloride  very  easily  becomes  so  modified  in 
solvent  power  by  hydrogen  chloride  dissolved  in  it  to  saturation 
that  barium  chloride  will  dissolve  in  the  solution  only  in  the 
ratio- of  1  to  about  20,000  parts.  The  addition  of  ethyl  ether 
to  the  concentrated  aqueous  hydrochloric  acid  thus  produced 
further  diminishes  the  solvent  action  of  the  mixture,  so  that 
in  a  mixture  of  concentrated  aqueous  hydrochloric  acid  with 
20  per  cent  of  its  volume  of  absolute  ether  the  salt  dissolves  in 
the  solvent  mixture  in  the  proportion  of  1  :  120,000  parts.  Ad- 
vantage may  be  taken  of  this  fact  to  precipitate  barium  chloride 
and  separate  it  from  calcium  chloride  and  magnesium  chloride 
which  remain  soluble  in  high  degree.  The  chlorides  of  these 
elements  are  dissolved  in  the  least  possible  amount  of  hot 
water  and,  after  cooling,  the  mixture  of  concentrated  hydro- 
chloric acid  and  ether  (5:1)  is  added  drop  by  drop  at  the  outset, 
to  induce  the  formation  of  the  precipitate  in  coarsely  crystalline 
condition.  After  standing  a  few  minutes,  the  precipitated 
hydrous  chloride,  BaCl2  •  2  H2O,  is  collected  upon  asbestos  in 
the  perforated  crucible,  washed  with  the  hydrochloric  acid- 
ether  mixture,  dried  at  150  to  200  degrees,  and  weighed  as  the 
anhydrous  chloride  BaCl2.* 

Instead  of  charging  the  water  solution  with  gaseous  hydro- 
chloric acid,  that  reagent  may  be  produced  by  the  action  of 
acetyl  chloride  upon  the  water,  inconvenient  violence  of  reaction 
being  moderated  by  the  addition  of  acetone  which  mixes  in  all 
proportions  with  both  acetyl  chloride  and  water: 

CHsCOCl  +  H20  =  CH3COOH  +  HC1. 

When  a  mixture  of  acetone  and  acetyl  chloride,  preferably 
4  :  1,  is  added  slowly  to  a  very  concentrated  solution  of  barium 
chloride  in  water,  the  water  is  attacked  at  once,  hydrogen 
chloride  is  liberated,  and  precipitation  begins  immediately.  If 
the  temperature  is  kept  down  during  the  process  by  immers- 
ing in  cool  running  water  the  vessel  in  which  reaction  takes 
place,  no  more  than  a  mere  trace  of  barium  can  be  detected 
by  testing  with  sulphuric  acid  the  residue  after  evaporating  the 
liquid  separated  from  the  precipitate  by  filtration  through 

*  Mar:  Methods  in  Chemical  Analysis,  Gooch,  p.  174. 


72  QUANTITATIVE  CHEMICAL  ANALYSIS 

asbestos.  When,  however,  the  temperature  is  allowed  to  rise, 
in  consequence  of  the  heat  liberated  in  the  reaction,  an  appre- 
ciable amount  of  barium  may  be  found  by  sulphuric  acid  in 
the  nitrate.  It  appears  that  when  the  acetone-acetyl  chloride 
mixture  (4  : 1)  acts  upon  the  cooled  concentrated  water  solu- 
tion of  barium  chloride  the  precipitate  is  the  hydrous  chloride, 
BaCl2  •  2  H2O,  while  the  water  in  excess  of  that  needed  to  form 
the  hydrous  salt  is  immediately  attacked;  that  acetyl  chloride 
by  itself  produces  only  slight  dehydration  of  this  salt  without 
marked  solubility;  and  that  prolonged  action  of  an  acetone- 
acetyl  chloride  mixture  (2  :  1)  results  in  an  appreciable  de- 
hydration and  considerably  increased  solubility  of  the  salt. 
When  the  acetone-acetyl  chloride  mixture  is  added  without 
cooling  to  the  water  solution  of  barium  chloride  the  heat  of 
reaction  favors  dehydration  of  the  hydrous  salt,  and  the  an- 
hydrous salt  may  go  into  solution  to  the  amount  of  several 
milligrams  in  10  cm.3  of  the  precipitating  mixture.  Upon 
filtering  the  mixture  and  treating  the  filtrate  with  acejone, 
with  acetyl  chloride,  or  with  the  acetone-acetyl  chloride  mix- 
ture, the  dissolved  anhydrous  salt  is  not  thrown  out  of  solution, 
but  the  addition  of  a  drop  of  water  is  sufficient  to  induce  im- 
mediate precipitation  in  the  form  of  the  hydrous  salt. 

The  best  conditions  for  the  quantitative  precipitation  of 
barium  chloride  by  the  acetone-acetyl  chloride  mixture  are 
found  in  the  use  of  minimum  amounts  of  water,-  the  preserva- 
tion of  ordinarily  low  temperature,  a  liberal  proportion  of  ace- 
tone, and  not  too  prolonged  digestion  of  the  precipitate  in  the 
excess  of  the  precipitant.  The  salt  to  be  analyzed  is  weighed 
out  into  a  small  beaker  and  dissolved  in  1  cm.3  of  water.  The 
beaker  is  cooled  by  immersion  in  a  water-bath,  preferably  sup- 
plied with  running  water,  at  a  temperature  of  about  15  degrees. 
To  the  copied  solution,  constantly  shaken,  the  acetone-acetyl 
chloride  mixture  is  added  from  a  dropping  funnel  at  the  rate  of 
five  drops  to  the  second.  The  precipitate  is  filtered  off  upon 
asbestos  in  a  perforated  crucible,  dried,  or  ignited,  and  weighed 
as  the  anhydrous  chloride,  BaCb  The  best  conditions  studied 
for  the  handling  of  0.1  grm.  of  hydrous  barium  chloride  are  the 
solution  of  the  salt  in  1  cm.3  of  water,  treatment  with  30  cm.3  of  the 
4 : 1  mixture  of  acetone  and  acetyl  chloride,  washing  with  acetone, 
and  drying  in  the  air-bath  at  135  degrees  or  at  low  redness.* 

Precipitation  of  Aluminium  Chloride.  —  Aluminium  chloride 
may  be  precipitated  similarly  by  thoroughly  saturating  the  cold 
aqueous  solution  of  the  chloride  with  hydrogen  chloride  after 

*  Gooch  and  Boynton:  Methods  in  Chemical  Analysis,  Gooch,  p.  175. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  73 

adding  an  equal  volume  of  ether.  The  precipitated  hydrous 
chloride,  AlCls  •  6  H2O,  undergoes  hydrolysis  when  carefully 
dried  and  ignited,  leaving  aluminium  oxide.  Any  slight  danger 
of  mechanical  loss  during  the  ignition  may  be  counteracted  by 
covering  the  precipitate  in  the  filtering  crucible  with  a  layer  of 
mercuric  oxide  which,  after  acting  as  a  sieve  to  prevent  the 
escape  of  solid  particles  during  the  decomposition  of  the  hydrous 
chloride,  will  finally  be  entirely  volatilized.  By  this  process 
aluminium  may  be  separated  from  iron,*  and  from  glucinum, 
mercury,  zinc,  copper,  and  bismuth.f 

In  the  practical  execution  of  the  process  the  concentrated 
aqueous  solution  of  the  salts  is  mixed  with  a  suitable  volume  of 
strongest  aqueous  hydrochloric  acid  —  enough  to  make  the  en- 
tire volume  approximately  15  cm.3  to  25  cm.3.  This  mixture 
is  saturated  with  gaseous  hydrochloric  acid  while  cooled  by  im- 
mersing the  container  in  running  water.  A  volume  of  ether 
equal  to  the  volume  of  the  liquid  is  introduced,  and  the  cooled 
etherial  mixture  is  treated  with  gaseous  hydrochloric  acid  to 
saturation  at  a  temperature  not  exceeding  15  degrees. J  The 
precipitated  crystalline  chloride  is  collected  upon  asbestos  in  a 
perforated  crucible,  washed  with  a  previously  prepared  mixture 
of  hydrochloric  acid  and  ether  carefully  saturated  at  15  degrees, 
dried  half  an  hour  at  150  degrees,  covered  with  a  layer  of  pure 
mercury  oxide  (about  1  grm.),  and  ignited  (under  a  ventilating 
flue)  gently  at  first  and  finally  with  the  blast-lamp. 

The  Precipitation  of  Sulphates 

The  Determination  of  the  Sulphate  Ion.  —  The  gravimetric 
determination  of  the  sulphate  ion  in  free  sulphuric  acid  and  in 
sulphates  rests  fundamentally  upon  the  reaction  which  may 
take  place  between  the  free  acid,  or  a  metal  sulphate,  and 
barium  chloride,  according  to  the  typical  reactions 
H2S04  +  BaCl2  <=±  BaS04  +  2  HC1, 
M2SO4  +  BaCl2  <=±  BaSO4  +  2  MCI. 

*  Gooch  and  Havens:   Methods  in  Chemical  Analysis,  Gooch,  p.  214. 

f  Havens:  Ibid.,  p.  216. 

|  The  gaseous  hydrochloric  acid  is  most  conveniently  produced  in  regu- 
lated current  by  treating  massive  ammonium  chloride  with  strong  sul- 
phuric acid  in  the  Kipp  generator.  A  platinum  dish  hung  in  an  inverted 
bell  jar,  provided  with  inlet  and  outlet  tubes  through  which  the  current 
of  water  for  cooling  is  passed,  makes  a  convenient  container  for  the  solu- 
tion to  be  saturated  with  the  gas. 


74  QUANTITATIVE  CHEMICAL  ANALYSIS 

In  regard  to  the  former  of  these  reactions  it  has  been  pointed 
out  *  that,  while  the  reaction  is  reversible,  the  concentration  of 
the  barium  sulphate  in  solution  is  always  very  low  (0.0023  grm. 
in  1  liter  of  pure  water)  and  is  only  slightly  affected  by  small 
changes  in  the  concentration  of  the  hydrochloric  acid.  The 
second  reaction,  that  between  a  soluble  sulphate  and  barium 
chloride,  is  likewise  reversible,  and  in  it  the  concentration  of 
dissolved  barium  sulphate  is  also  very  low  and  not  very  much 
affected  by  changes  in  the  concentration  of  the  chloride  pro- 
duced. The  determination  of  the  sulphuric  acid  ion  should, 
therefore,  be  a  comparatively  simple  matter  if  the  reactions  in- 
dicated were  the  only  possible  reactions  between  barium  chlo- 
ride and  sulphuric  acid  or  a  sulphate.  The  fact  is,  however, 
that  other  reactions  intervene. 

Contamination  of  Precipitated  Barium  Sulphate.  —  Wheji  a 
solution  of  barium  chloride  is  added  to  a  dilute  solution  of 
.sulphuric  acid  (kept  at  the  boiling  temperature  in  order  that 
the  precipitate  of  barium  sulphate  may  be  coarse  enough  to 
settle  well)  very  nearly  all  of  the  sulphuric  acid  ion  is  precipi- 
tated, but  the  barium  sulphate  includes  some  barium  chloride, 
not  removable  by  washing;  and  the  inclusion  is  greater  when 
the  barium  chloride  is  added  rapidly.  Furthermore,  it  has  been 
found  that  the  precipitate  may  yield  upon  ignition  sulphuric 
acid  or  hydrochloric  acid,  or  both.  When  barium  reacts  with 
a  soluble  sulphate  still  other  complexities  may  come  to  light. 
The  precipitated  sulphate  may  not  only  include  barium  chloride 
but  it  may  contain  other  sulphates.  When,  for  example,  barium 
chloride  and  potassium  sulphate  react  in  solution  the  precipitate 
may  contain  potassium  sulphate  amounting  to  as  much  as 
several  per  cent  of  the  total  weight. 

To  explain  the  obstinate  retention  of  barium  chloride  in  the 
precipitated  barium  sulphate  and  the  evolution  of  the  free 
acids  upon  ignition  it  has  been  suggested  f  that  combinations, 
such  as  Bad  •  HS04  and  BaC!2  •  BaS04,  are  formed,  which  may 
be  figured  as  shown  in  the  following  symbols: 

*  Page  8. 

f  Hulett  and  Duschak:  Zeit.  anorg.  Chem.,  40,  196  (1904). 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  75 

Cl  Cl 


O  O' 


The  obstinate  retention  of  an  ordinarily  soluble  sulphate  by 
precipitated  barium  sulphate  may,  likewise,  be  explained  by 
assuming  the  formation  of  double  sulphates  derived  from 
sulphuric  acid  or  a  soluble  sulphate  existing  in  molecules  more 
complex  than  those  which  are  represented  by  the  simple  equiva- 
lent symbols.  If,  for  example,  some,  at  least,  of  the  molecules 
of  sulphuric  acid  and  of  the  sulphate  are  present  in  polymeric 
form,  it  is  easy  to  see  how  complex  molecules  may  result,  such 
as  BaS04  •  H2S04  and  BaS04  •  M2SO4,  which  may  be  figured  as 

/°\  ^  /°-H      /°\  /rN  /0-M 
Ba(  ><^>S(        Ba(  ><^>s( 

X0X  xgx  X0-H      ^O7  \g/  X0-M 

The  formation  of  such  complex  substances  must  introduce 
many  inaccuracies  in  determinations  based  upon  the  funda- 
mental presumption  that  barium  sulphate,  BaSO4,  is  the  final 
and  only  product  weighed.  Thus,  the  product  BaCl2  •  BaSO4 
will  weigh  too  much  by  the  amount  of  BaCl2  in  it;  BaCl  •  HSO4 
may  lose  on  ignition  hydrochloric  acid  and  become  normal  in 
composition  and  weight,  or  it  may  lose  sulphuric  acid  and  be 
deficient  in  weight;  BaS04  •  H2SO4  will  lose  sulphuric  acid  on 

ignition,  weighing  too  little  by  the  amount  of  loss;    and  the 

i 
weight  of  BaSO4  •  M2SO4  will  differ  from  what  the  weight  of 

the  barium  sulphate  should  be  by  an  amount  depending  upon 

the  difference  between  the  atomic  weights  of  Ba  and  M2. 

Precaution  in  Precipitations.  —  In  the  practical  determina- 
tion of  the  sulphate  ion  of  free  sulphuric  acid,  and  of  certain 
sulphates  (e.g.,  copper  sulphate,  magnesium  sulphate,  and  alkali 
sulphates  in  absence  of  a  large  excess  of  alkali  salts)  results 


76  QUANTITATIVE  CHEMICAL  ANALYSIS 

very  near  the  truth  may  be  obtained  by  adding,  drop  by  drop, 
to  the  hot  dilute  solution  of  the  acid  or  sulphate  a  dilute  solu- 
tion of  an  equivalent  amount  of  barium  chloride,  and  when  the 
precipitation  is  practically  complete  a  considerable  excess  of 
that  reagent  to  bring  about  the  highest  degree  of  insolubility  of 
the  barium  sulphate.  Under  such  circumstances  the  inclusion 
of  barium  chloride  is  at  a  minimum  and  the  increase  of  weight, 
due  to  such  inclusion,  is  approximately  equal  to  that  of  the 
barium  sulphate  lost  in  consequence  of  solubility. 

Experimental  Process:  the  Sulphate  Ion  in  Copper  Sulphate.  —  Weigh 
out  carefully  about  0.5  grm.  of  copper  sulphate,  CuSO4  •  5  H2O,  transfer 
it  to  a  beaker,  dissolve  it  in  about  400  cm.3  of  water,  heat  the  solution  to 
boiling,  and  add  0.5  cm.3  of  concentrated  hydrochloric  acid.  To  the  solu- 
tion, at  the  boiling  temperature,  add  slowly  from  a  dropping  funnel  or 
burette,  with  constant  stirring,  25  cm.3  of  a  solution  of  barium  chloride, 
BaCl2  •  2  H2O,  containing  25  grm.  of  this  salt  to  the  liter.  This  amount 
of  the  precipitant  is  approximately  that  required  to  bring  about  the  re- 
action 

BaCl2  +  CuS04  <=*  BaS04  +  CuCl2 

Then  add  25  cm.3  more  of  the  precipitant  and  let  the  precipitate  settle 
for  half  an  hour. 

A.  Filtration  on  Paper. — Using  a  rod  held  vertically  to  direct  the  flow, 
decant  upon  the  filter  the  clear  liquid,  above  the  settled  precipitate,  to  the 

lowest  possible  limit,  taking  care  in  filling 
the  filter  to  leave  a  margin  half  a  centi- 
meter broad  above  the  level  of  the  liquid. 
Fill  the  beaker  with  distilled  water  to  the 
original  level  of  the  liquid,  settle  the 
precipitate  and  decant  as  before.  Finally 
transfer  the  precipitate  to  the  filter,  wash- 
ing out  the  beaker  carefully  and  dislodg- 
FIG.  19.  ing  any  adherent  precipitate  with  the  aid 

of  a  rubber-tipped  glass  rod  (the  "police- 
man")- Wash  the  precipitate  into  the  point  of  the  filter  with  a  fine  jet 
of  water  from  the  wash-bottle,  directing  the  jet  with  a  rotary  motion 
around  the  paper  from  the  rim  downward,  and  continue  the  washing  until 
silver  nitrate  fails  to  show  the  presence  of  a  chloride  in  a  portion  of  the 
washings  collected  separately.  Dry  the  filter  in  the  funnel,  at  least 
partially,  fold  it  compactly  in  a  weighed  platinum  crucible,  cover  the 
crucible,  and  heat  very  gently  until  dry.  Incline  the  crucible  (Fig.  19), 
arrange  the  cover  so  as  to  make  a  draft  of  hot  air  from  the  flame  through 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


77 


it,  increase  the  heat  until  the  paper  begins  to  char,  but  keeping  it  low 
until  all  volatile  matter  has  escaped.  Then  move  the  flame  so  that  it 
strikes  the  bottom  of  the  inclined  crucible  and  increase  the  heat  to  redness 
until  all  carbon  has  been  consumed.  It  is  important  not  to  raise  the  tem- 
perature to  too  high  a  point  in  the  early  stages  of  the  ignition  lest  the 
barium  sulphate  be  reduced  and  sulphur  liberated  appreciably,  and  not  to 
heat  finally  in  the  blast  lamp  lest  sulphur  trioxide  be  expelled. 

Cool  the  crucible  nearly  to  the  room  temperature,  place  it  in  the  desic- 
cator and  keep  it  there  (ten  minutes)  until  thoroughly  cool,  and  weigh. 
Moisten  the  residue  with  sulphuric  acid,  heat  gently  within  another  crucible 
or  in  an  iron  cone,  used  as  a  radiator,  and  ignite  to  a  constant  weight. 
Take  the  difference  between  the  weight  thus  found  and  the  weight  of  the 
crucible  as  that  of  the  barium  sulphate.  Calculate  the  weight  of  the  SO4 
ion  in  the  barium  sulphate  and  compare  it  with  that  which  should  be 
obtained  from  the  copper  sulphate  taken.  State  the  results  in  a  scheme 
like  the  following: 


CuSO4  •  5  H2O 
taken. 

BaS04 
found. 

SO4 
calculated. 

Theory  for  SO4  in 
CuSO4  •  5  H2O  taken. 

Error  in  the 
determination  of  SO4. 

Grm. 

Grm. 

Grm. 

Grm. 

Per  cent. 

Grm. 

Per  cent. 

B.  Filtration  on  Asbestos.  — Place  in  the  rubber  holder  of  a  vacuum  flask 
a  perforated  crucible,  fitted  with  an  asbestos  felt  (as  previously  described, 
p.  64),  ignited,  and  weighed.  Start  the  pump,  moisten  the  asbestos  with 
a  little  water,  and,  after  the  crucible  has  settled  into  place,  decant  into  it 
the  clear  supernatant  liquid.  Transfer  the  precipitate  to  the  crucible 
and  wash  it,  directing  the  jet  from  the  wash-bottle  upon  the  wall  of  the 
crucible,  in  order  to  avoid  any  disturbance  of  the  asbestos.  Place  the 
crucible  containing  the  asbestos  felt  and  the  precipitate  upon  a  pipe-stem 
triangle  of  such  size  that  the  cap*  which  covers  the  bottom  will  rest 
firmly.  Heat  with  a  Bunsen  burner,  very  gently  at  first,  so  that  the  felt 
will  not  be  stirred  by  escaping  steam,  then  to  red  heat,  cool  in  the  desiccator 
and  weigh.  Arrange  the  results  of  the  analysis  in  the  form  given  for  the 
preceding  process. 

General  Procedure.  —  Beside  the  contamination  of  the  pre- 
cipitated barium   sulphate  there  is  sometimes   another   diffi- 
culty, as  when  salts  of  chromium  or  iron  are  present,  due  to 
the  fact  that  the  precipitation  of  some  of  the  sulphate  ion  is 
*  Blair's  pattern. 


78  QUANTITATIVE  CHEMICAL  ANALYSIS 

retarded  or  prevented,  probably  by  the  formation  of  complex 
ions.  Generally,  therefore,  it  is  better  in  determining  the 
sulphate  ion  of  sulphates,  and  in  many  cases  necessary,  to  re- 
move the  basic  element  by  a  preliminary  precipitation,  before 
attempting  the  precipitation  of  the  sulphate  ion  by  the  soluble 
'barium  salt;  although  the  compounds  of  certain  bivalent 
elements,  like  magnesium  and  copper,  when  present  in  small 
amount,  do  not  interfere  seriously. 

In  practical  analysis  it  is  often  necessary  to  determine  the 
sulphuric  acid  ion  in  the  presence  of  considerable  amounts  of 
alkali  salts.  For  example,  in  certain  processes,  such  as  the  de- 
termination of  sulphur  in  pyrites,  in  which  it  is  usual  to  oxidize 
the  sulphur  by  fusion  with  sodium  carbonate  mixed  with 
sodium  peroxide,  or  other  oxidizer  (protected  from  the  action 
of  sulphur  compounds  in  the  fuel  gases*),  sodium  sulphate  is 
obtained  in. a  mass  of  sodium  carbonate  with  insoluble  oxides 
or  carbonates.  Upon  treatment  with  water  this  mass  yields  the 
sulphate  in  solution  with  sodium  carbonate  and,  perhaps,  sodium 
silicate.  By  acidifying  the  solution  with  hydrochloric  acid  and 
evaporating  to  dryness,  silica  is  made  nearly  insoluble  f  and, 
upon  extracting  with  water  containing  hydrochloric  acid  and 
filtering,  it  is  nearly  completely  removed,  while  the  solution 
holds  the  sodium  sulphate  with  much  sodium  chloride.  To  get 
the  best  results  the  larger  part  of  the  sodium  chloride  should  be 
removed  before  attempting  the  precipitation  of  the  sulphate  by 
means  of  barium  chloride;  and  this  may  be  accomplished  by 
adding  to  the  solution,  previously  concentrated  to  the  lowest 
limit,  five  or  six  times  its  volume  of  concentrated  hydrochloric 
acid,  filtering  on  asbestos  in  the  perforated  crucible,  and  washing 
the  residue  of  sodium  chloride  with  concentrated  hydrochloric 
acid.  The  soluble  sulphate  is  thus  separated  from  the  greater 
part  of  the  sodium  chloride,  insoluble  to  a  very  considerable  ex- 
tent in  the  concentrated  hydrochloric  acid.  After  diluting  (to 

*  The  platinum  crucible  in  which  the  fusion  is  made  may  be  fitted  into 
a  circular  hole  cut  through  the  middle  of  an  asbestos  disk  and  edged  with 
platinum  foil,  the  heating  flame  being  applied  from  below. 

t  See  p.  93. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  79 

prevent  loss  during  the  evaporation)  and  evaporating  (to  re- 
move the  free  hydrochloric  acid)  the  sulphate  ion  may  be 
determined  by  precipitation  in  the  form  of  reasonably  pure 
barium  sulphate.* 

The  Determination  of  Barium.  —  In  precipitating  barium 
sulphate  by  the  action  of  an  excess  of  sulphuric  acid  upon 
barium  chloride  in  solution  the  barium  is  held  in  the  precipi- 
tate (excepting  the  extremely  small  amount  which  dissolves  as 
sulphate)  but  not  all  in  the  form  of  the  sulphate.  If  the  precipi- 
tate is  weighed  the  result  is  low  upon  the  hypothesis  that 
barium  sulphate  is  the  sole  product  in  which  barium  appears; 
moreover,  chlorine  may  be  detected  in  the  precipitate.  As  in 
the  reverse  precipitation  of  sulphuric  acid  by  barium  chloride, 
the  precipitate  contains  barium,  the  sulphate  ion,  .and  chlorine; 
but  in  the  present  case  it  is  possible  to  make  the  necessary  cor- 
rection by  simply  moistening  the  precipitate  with  sulphuric 
acid  and  volatilizing  by  gentle  ignition  the  excess  with  the  hy- 
drochloric acid  produced  in  the  reaction: 

BaCl2  •  BaS04  +  H2S04  =  2  BaS04  +  HC1. 

In  the  precipitation  of  barium  by  sulphuric  acid  in  the  pres- 
ence of  alkali  salts,  as  in  the  process  in  which  barium  sulphate 
is  precipitated  by  the  action  of  barium  chloride  upon  alkali  sul- 
phates, the  precipitate  is  in  like  manner  contaminated  by  alkali 
sulphates  —  sometimes  as  much  as  20  mgrm.  in  0.5  .grm.  of  pre- 
cipitate. In  this  case,  however,  advantage  may  be  taken  of  the 
fact  that  barium  sulphate  may  be  dissolved  in  concentrated 
sulphuric  acid  and,  by  evaporation  of  the  acid,  recovered  in 
crystalline  grains  from  which  the  alkali  sulphates  may  be  washed 
with  water.  Fig.  20  shows  a  simple  devicef  for  effecting  the 
purification  of  barium  sulphate  from  alkali  sulphates,  by  solu- 
tion in  sulphuric  acid,  recrystallization  and  evaporation  of  the 
acid.  The  precipitate,  filtered  upon  ashless  paper,  is  ignited 
with  the  paper  in  a  weighed  platinum  crucible  and  concentrated 
sulphuric  acid  (2  cm.3  to  5  cm.3)  is  added.  Within  the  mouth 

*  Turner:   Am.  Jour.  Sci.,  [4],  38,  41. 

t  Gooch  and  Hill:   Am.  Jour.  Sci.,  35,  311  (1913). 


80  QUANTITATIVE  CHEMICAL  ANALYSIS 

of  the  crucible  a  weighed  cone  of  fine  platinum  gauze  is  fitted 
to  prevent  loss  by  spattering,  and  the  flame  of  blast-lamp  is 
directed  into  the  point  of  the  cone  from  above.  When  the 
evaporation  of  the  excess  of  sulphuric  acid  is  completed,  the  resi- 
due is  washed  with  water  and  that  portion  of  the  barium  sul- 
phate which  is  removed  from 
the  crucible  in  this  process 
is  collected  upon  a  filter  of 
ashless  paper.  The  filter  is 
dried  and  ignited  in  the 
crucible.  The  crucible,  cone, 
and  the  residue  of  purified 
barium  sulphate  are  weighed 
together. 

Experimental    Process :     Deter- 
mination  of  Barium   as    Barium 

-pIG   20  Sulphate.  — Dissolve  in  hot  water 

about  0.5  grm.  of  barium  chloride, 
BaCl2  •  2  H2O,  weighed  exactly,  dilute  the  solution  to  a  volume  of  about 
400  cm.3,  heat  it  to  boiling,  add  dilute  [1:4]  sulphuric  acid  in  excess,  boil 
a  few  minutes,. settle,  filter  on  asbestos  in  the  perforated  crucible,  wash, 
dry,  ignite,  and  weigh.  Moisten  the  residue  with  a  drop  of  dilute 
sulphuric  acid,  ignite,  and  weigh  as  BaSO4.  This  process  will  give  good 
results  in  presence  of  large  amounts  of  hydrochloric  acid,  provided  the 
excess  of  sulphuric  acid  is  large,  e.g.,  1  to  2  per  cent  by  volume  of 
the  entire  solution. 

The  Determination  of  Strontium.  —  Strontium  may  be 
precipitated  and  determined  as  strontium  sulphate  provided 
means  be  used  to  properly  decrease  the  solubility  of  the  pre- 
cipitate—  about  1  :  7000  in  water  at  the  atmospheric  tem- 
perature and  somewhat  more  at  the  boiling  temperature,  and 
considerably  more  in  presence  of  hydrochloric  acid,  nitric  acid, 
or  salts  of  these  acids. 

SrCl2  +  H2S04  <=*  SrS04  +  2  HC1. 

The  addition  of  an  equal  volume  of  alcohol  to  the  solution  after 
precipitation  by  a  slight  excess  of  sulphuric  acid  will  render 
the  insolubility  sufficient  for  analytical  purposes.  The  precipi- 
tate, after  standing  twelve  hours,  is  collected  upon  asbestos  in 
the  perforated  crucible,  washed  with  50  per  cent  alcohol  con- 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  81 

taining  a  little  sulphuric  acid,  dried,  and  ignited  at  low  red 
heat. 

The  Determination  of  Calcium.  —  Calcium  may  be  pre- 
cipitated as  calcium  sulphate  by  adding  an  excess  of  sulphuric 
acid  to  the  solution  of  the  calcium  salt  and  a  volume  of  alcohol 
equal  to  four  times  that  of  the  aqueous  solution  and  allowing 
the  mixture  to  stand  for  twelve  hours. 

CaCl2  +  H2S04  +±  CaS04  -f  2  HC1. 

The  precipitate  is  collected  on  asbestos  in  the  perforated 
crucible,  dried,  and  ignited  at  low  red  heat. 

The  Determination  of  Lead.  —  From  the  water  solution  of 
lead  chloride,  lead  nitrate,  or  lead  acetate,  the  lead  may  be 
precipitated  as  lead  sulphate  by  the  addition  of  an  excess  of 
sulphuric  acid  and  a  volume  of  alcohol  twice  that  of  the  water 

solution. 

Pb(N03)2  +  H2S04  <±  PbS04  +  2  HNO3. 

If  the  addition  of  alcohol  is  undesirable  on  account  of  sub- 
sequent analytical  operations  the  precipitate  may  be  made 
nearly  insoluble  by  the  addition  of  a  large  excess  of  sulphuric 
acid  (e.g.,  10  per  cent  by  volume  of  the  entire  solution). 

Experimental  Process:  Determination  of  Lead  as  Lead  Sulphate.  — 
A.  Weigh  out  carefully  about  0.5  grm.  of  powdered  and  dried  lead  nitrate, 
dissolve  it  in  water  (50  cm.3)  with  the  aid  of  a  drop  or  two  of  nitric  acid 
(to  dissolve  a  trace  of  basic  salt),  and  add  dilute  sulphuric  acid,  in  excess, 
with  a  volume  of  alcohol  twice  as  great  as  that  of  the  water  solution. 
Settle  the  precipitate,  filter  it  off  on  asbestos  in  the  perforated  crucible, 
wash  with  alcohol  (80  per  cent),  ignite,  and  weigh  as  PbSO4. 

B.  Take  about  0.5  grm.  of  powdered  and  dried  lead  nitrate  weighed 
exactly.  Dissolve  it  in  water  (50  cm.3)  with  the  aid  of  a  drop  or  two  of 
nitric  acid.  Add  a  volume  of  dilute  sulphuric  acid  (1:4)  equal  to  that 
of  the  water  solution.  Filter  on  asbestos  in  the  perforated  crucible,  wash 
with  dilute  sulphuric  acid  (1  :  4),  ignite,  and  weigh  as  PbSO4. 

The  Precipitation  of  Phosphates 
The  Determination  of  the  Orthophosphate  Ion.  —  One  of 

the  processes  by  which  the  acidic  ion  of  free  orthophosphoric 
acid  and  its  salts  may  be  determined  depends  upon  the  re- 
action in  which  ammonium  magnesium  phosphate,  NH4MgPC>4  • 
6  H20,  is  precipitated  from  solution  by  the  action  of  the  soluble 


82  QUANTITATIVE  CHEMICAL  ANALYSIS 

"  magnesia  mixture,"  made  up  by  putting  together  in  solution 
magnesium  chloride,  ammonium  chloride,  and  ammonium 
hydroxide.  The  ideal  reaction  proceeds  in  the  sense  of  the 
expression 

MgCl2  +  NH4OH  +  HNa2P04  +  nNH4Cl  =  NH4MgP04 
+  2  NaCl  +  H20  + 


The  function  of  the  ammonium  chloride  is  the  prevention  of 
the  precipitation  of  magnesium  hydroxide  so  that  the  compo- 
sition of  the  precipitate  may  be  definite  and  ideal.  From  the 
precipitate,  submitted  to  ignition,  ammonia  and  water  are 
eliminated  while  magnesium  pyrophosphate  remains  in  weigh- 
able  form. 


NH4-0X 

0 


Mg(        P=0 

X0X 

As  little  as  0.0001  grm.  of  the  orthophosphoric  acid  ion  may  be 
precipitated  in  500  cm.3  of  faintly  ammoniacal  solution,  even 
in  presence  of  as  much  as  60  grm.  of  ammonium  chloride,  by  a 
sufficient  excess  of  the  magnesium  chloride;  but  the  ideal 
ammonium  magnesium  phosphate,  NH4MgP04,  which  yields 
upon  ignition  the  pyrophosphate,  Mg2P2O7,  may,  according  to 
conditions  of  precipitation,  be  contaminated  by  the  trimag- 
nesic  phosphate  which  upon  ignition  gives  the  anhydrous 
phosphate  Mg3P2Os,  or  by  a  double  phosphate,  (NH4)4Mg(P04)2 
which  upon  ignition  leaves  magnesium  metaphosphate  Mg(P03)2. 
The  ammonium  salts  tend  to  bring  about  the  contamination 
of  the  precipitate  which  results  in  the  formation  of  the  meta- 
phosphate on  ignition,  and  the  use  of  strongly  ammoniacal 
solutions  induces  contamination  by  the  trimagnesic  phosphate. 
The  results  of  practical  experiments  go  to  show  that  the  con- 
stitution of  the  precipitate  is  very  nearly  ideal  when  the 
boiling  solution  of  the  phosphate,  containing  a  moderate  ex- 
cess of  the  magnesium  salt  and  not  more  than  5  to  10  per  cent 
of  ammonium  chloride  is  made  ammoniacal  very  gradually,  so 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  83 

that  the  precipitate  is  formed  in  crystalline  condition.  If 
more  ammonium  chloride  than  this  proportion,  or  more  mag- 
nesium salt  than  twice  the  amount  theoretically  necessary,  is 
present,  it  is  safer  to  decant  the  supernatant  liquid  from  the 
precipitate  (through  the  filter  to  be  used  subsequently  to  hold 
the  phosphate),  to  dissolve  the  precipitate  in  a  little  hydro- 
chloric acid  and  then  to  reprecipitate  by  dilute  ammonia. 

The  Determination  of  Magnesium.  —  As  little  as  0.0001 
grm.  of  magnesium  may  be  precipitated  in  500  cm.3  of  a  faintly 
ammoniacal  solution,  even  when  the  solution  contains  as 
much  as  60  grm.  of  ammonium  chloride  (or  100  cm.3  of  a 
saturated  solution  of  ammonium  oxalate).  As  in  the  precipi- 
tation of  ammonium  magnesium  phosphate  in  the  determina- 
tion of  the  orthophosphoric  acid  ion,  much  ammonia  tends  to 
the  formation  of  the  trimagnesium  phosphate  while  a  large 
amount  of  ammoniun  salts  tend  to  introduce  too  much  am- 
monium into  the  double  phosphate.* 

When  ammonium  salts  are  present  in  quantity,  as  is  often  the 
case  in  the  ordinary  course  of  analysis,  the  precipitate  first 
thrown  down  by  addition  of  ammonium  sodium  phosphate  and 
ammonia  in  distinct  excess  should  be  settled  and  the  super- 
natant liquid  poured  off  through  the  filter  used  subsequently 
in  collecting  the  precipitate.  The  precipitate  is  then  dissolved 
in  the  least  possible  amount  of  hydrochloric  acid  and  thrown 
down  again  from  the  diluted  solution  by  ammonia  in  slight 
excess.  For  safety,  a  little  ammonium  sodium  phosphate  may 
also  be  added.  The  precipitate,  filtered  off  and  washed  with 
distinctly  ammoniacal  water,  is  ignited  very  gently  at  first 
(to  avoid  reduction  of  the  phosphoric  oxide  by  the  ammonia 
set  free)  until  all  ammonia  is  expelled,  when  the  temperature 
is  raised  to  redness. 

Good  results  are  also  obtained  (GMs)  by  adding  to  the 
boiling  solution  of  the  magnesium  salt  a  solution  of  ammonium 
sodium  phosphate  until  no  further  precipitation  takes  place 
and  then  (dilute)  ammonium  hydroxide  in  excess.  Even  in 
presence  of  considerable  amounts  of  ammonium  chloride  this 

*  See  p.  82. 


84  QUANTITATIVE  CHEMICAL  ANALYSIS 

process  yields  a  phosphate  of  nearly  ideal  constitution  if  only 
the  boiling  be  prolonged  from  three  to  five  minutes.  The 
greater  part  of  the  ammonium  magnesium  phosphate  —  about 
90  per  cent  —  forms  in  this  process  before  free  ammonia  is 
added,  arid  the  ammonium  which  enters  the  phosphate  thus 
formed  is  derived  from  the  ammonium  sodium  phosphate, 
which  becomes  correspondingly  acidic.  Under  these  conditions, 
the  tendency  to  form  an  insoluble  ammonium  magnesium  phos- 
phate, richer  in  ammonium  and  poorer  in  magnesia  than  the 
normal  salt,  is  slight. 

Experimental  Process:  Determination  of  Magnesium  as  Magnesium 
Pyrophosphate.  —  I.  Draw  from  a  burette  50  cm.3  of  a  solution  of  magne- 
sium chloride  (about  N/10  in  respect  to  the  anhydrous  salt)  standardized 
as  described  in  the  Experimental  Process  of  page  48.  Dilute  the  solution 
to  a  volume  of  150  cm.3  Add  in  solution  2.5  grm.  of  microcosmic  salt, 
4  Na(NH4)PO4  •  4  H2O,  dissolve  the  precipitate  which  comes  dbwn  in 
flocculent  condition  in  hydrochloric  acid  and  reprecipitate  in  crystalline 
condition  by  the  gradual  addition  of  ammonium  hydroxide  in  distinct 
excess.  Settle  the  precipitate,  pour  off  the  supernatant  liquid  through 
asbestos  weighed  in  the  perforated  crucible.  Dissolve  the  precipitate  in 
the  least  possible  amount  of  hydrochloric  acid,  dilute  the  solution  to  a 
volume  of  100  cm.3  or  150  cm.3,  reprecipitate  by  the  very  gradual  addition 
of  ammonium  hydroxide  in  excess,  filter  on  the  asbestos  previously  weighed 
and  used  in  the  decantation,  and  wash  with  distinctly  ammoniacal  water. 
Heat  the  precipitate  gently  at  first,  until  all  ammonia  has  been  expelled 
(to  avoid  reducing  action),  and  afterward  at  full  red  heat.  The  "glow" 
which  passes  over  the  gently  ignited  residue  when  heated  to  redness, 
marking  a  molecular  change  in  the  pyrophosphate,  is  a  good  indication 
that  the  ignition  has  been  sufficiently  prolonged.  Weigh  as  Mg2P2O7. 

II.  Draw  from  a  burette  5  cm.3  of  the  standardized  solution  of  mag- 
nesium chloride.  Add  ammonium  hydroxide  in  faint  excess  and  redis- 
solve  any  precipitated  magnesium  hydroxide  by  the  cautious  addition  of 
hydrochloric  acid.  To  the  neutral  or  very  slightly  acid  solution  heated 
to  boiling  add  gradually  a  considerable  excess  of  microcosmic  salt, 
HNa(NH4)PO4  •  4  H2O,  boil  for  five  minutes,  then  cool,  add  an  excess  of 
ammonium  hydroxide  and  set  aside  to  allow  the  precipitate  to  become 
crystalline.  Filter  on  asbestos  in  the  perforated  crucible,  wash  with  dis- 
tinctly ammoniacal  water,  ignite  gently  at  first  and  then  at  full  red  heat. 
Weigh  as  Mg2P2O7. 

The  Determination  of  Manganese.  —  The  manganese  phos- 
phate when  precipitated  in  the  cold  by  an  excess  of  an  alkali 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  85 

phosphate  in  the  presence  of  ammonium  chloride  and  ammonia 
is  very  largely  in  the  form  of  the  trimanganese  phosphate, 

3  MnCl2  +  2  NH4OH  +  2  H(NH4)NaP04 

=  Mn3P2O8  +  2  NaCl  +  4  NH4C1  +  2  H2O, 

but  boiling  or  even  subsequent  standing  may  effect  a  more  or 
less  complete  conversion  of  the  manganese  phosphate  to  the 
ammonium  manganese  phosphate.  The  success  of  the  analyti- 
cal process  in  which  manganese  is  weighed  as  the  pyrophosphate 
(Gibbs)  turns,  therefore,  upon  the  change  of  the  trimanganese 
phosphate,  Mn3P208,  to  the  ammonium  manganese  phosphate, 
NH4MnP04;  and  the  presence  of  a  large  amount  of  ammonium 
salt  is  essential  to  the  formation  of  the  precipitate  of  ideal 
constitution.*  The  proportion  of  ammonium  chloride  present 
to  ammonium  manganese  phosphate  formed  should  be  at  least 
40  :  1,  or,  speaking  approximately,  200  molecules  of  ammonium 
chloride  must  be  present  in  the  liquid  to  every  molecule  of 
the  phosphate  formed;  and  the  ammonium  chloride  may  be  in- 
creased almost  to  the  point  of  saturation  of  the  liquid  without 
causing  more  than  a  trifling  solubility  of  the  ammonium  man- 
ganese phosphate  in  the  presence  of  an  excess  of  the  precipitant. 
Furthermore,  the  precipitate  may  be  washed  with  perfect  safety 
with  pure  water  as  well  as  with  slightly  ammoniacal  water. 
The  finely  granular  precipitate  which  may  be  obtained  by  slow 
action  of  dilute  ammonia  added  gradually  to  the  hot  solution 
of  the  manganese  salt  apparently  includes  a  portion  of  uncon- 
verted phosphate  which  resists  the  replacement  of  the  man- 
ganese by  ammonium.  On  the  other  hand,  the  precipitate  of 
flocky  condition  thrown  down  in  the  cold  passes  easily  to  the 
silky  and  crystalline  condition  when  heated  with  the  proper 
amount  of  ammonium  salt,  and  possesses  a  constitution  ap- 
proaching the  ideal.  In  the  determination  of  manganese  by 
this  method  the  presence  of  a  large  amount  of  ammonium  chlo- 
ride is  therefore  necessary.  Good  results  may  be  obtained 
most  easily  and  surely  by  the  following  procedure:  The 
slightly  acid  solution,  containing  in  a  volume  of  200  cm.3 
*  Gooch  and  Austin:  Methods  in  Chemical  Analysis,  Gooch,  p.  482. 


86  QUANTITATIVE  CHEMICAL  ANALYSIS 

an  amount  of  manganese  not  more  than  enough  to  make  0.4 
grm.  of  the  pyrophosphate,  20  grm.  of  ammonium  chloride,  and 
5  cm.3  to  10  cm.3  of  a  cold  saturated  solution  of  ammonium 
sodium  phosphate,  is  precipitated  in  the  cold  by  the  careful 
addition  of  dilute  ammonia  in  slight  excess  only.  The  mixture 
is  heated  until  the  precipitate  becomes  silky  and  crystalline; 
the  whole  is  allowed  to  stand  and  cool  half  an  hour;  the  precipi- 
tate is  collected  upon  asbestos  in  a  perforated  platinum  crucible, 
washed  (best  with  slightly  ammoniacal  water),  dried  at  gentle 
heat,  and  ignited  as  usual,  in  the  presence  of  large  amounts  of 
ammonia. 

Experimental  Process:  Determination  of  Manganese  as  Manganous 
Pyrophosphate,  Mn^P^Oi.  —  Put  into  a  weighed  crucible  about  1  grm.  of 
hydrous  manganese  sulphate,  dehydrate  over  a  radiator  (according  to  the 
process  described  on  p.  49)  to  constant  weight  and,  noting  the  weight, 
dissolve  the  anhydrous  sulphate  in  water  containing  hydrochloric  acid. 
Add  20  grm.  of  purified  ammonium  chloride,*  5  cm.3  to  6  cm.3  of  a  satu- 
rated solution  of  hydrogen  sodium  ammonium  phosphate,  HNa(NH4)PO4, 
and  ammonium  hydroxide  drop  by  drop  in  slight  excess.  Raise  the  tem- 
perature of  the  solution  to  the  boiling  point  and  keep  it  there  until  the 
precipitate  becomes  crystalline.  Cool,  filter  on  asbestos  in  the  perforated 
crucible,  wash  with  distinctly  ammoniacal  water  (containing  at  the  end 
ammonium  nitrate).  Ignite,  gently  at  first  (to  avoid  reducing  action  of 
liberated  ammonia),  and  then  at  the  full  heat  of  the  Bunsen  burner. 
Weigh  as  Mn2P2O7. 

Cadmium  and  Zinc.  —  With  proper  precautions,  zinc  f  and 
cadmium }  may  be  likewise  precipitated  from  solutions  of  their 
soluble  salts  as  double  ammonium  phosphates  which  on  ig- 
nition leave  pyrophosphates. 

The  Precipitation  of  Oxalates 

The  Precipitation  of  Calcium  Oxalate.  —  Calcium  oxalate 
is  extremely  insoluble  in  water  and  in  ammoniacal  solutions, 
very  slightly  soluble  in  acetic  acid  or  oxalic  acid,  but  readily 

*  Prepared  by  boiling  the  solution  of  the  C.  P.  salt  of  commerce  with 
ammonium  hydroxide  and  filtering. 

t  Austin:  Methods  in  Analysis,  Gooch,  p.  185.  Treadwell-Hall : 
Analytical  Chemistry,  3rd  ed.,  Vol.  II,  p.  140. 

J  Austin:  Methods  in  Analysis,  Gooch,  p.  190. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  87 

dissolved  by  the  stronger  acids  in  fairly  concentrated  solutions. 
When  precipitated  from  cold  solutions  calcium  oxalate  is  so 
finely  divided  that  it  will  settle  only  very  slowly  and  on  filtra- 
tion tends  to  pass  through  the  pores  of  the  filter,  but  if  pre- 
cipitated from  boiling  hot  solution  and  allowed  to  stand  the 
precipitate  takes  a  more  coarsely  crystalline  form  and  is  then 
easily  filterable. 

From  neutral  or  slightly  ammoniacal  solutions  containing 
only  alkali  salts  beside  the  soluble  calcium  salt,  calcium  oxalate 
may  be  precipitated  in  reasonably  pure  condition  by  ammo- 
nium oxalate: 

CaCl2  +  2  (NH4)202C202  <=>  Ca02C202  +  2  NH4C1. 

But  if  a  magnesium  salt  is  also  present,  as  is  frequently  the  case 
in  practical  analysis,  magnesium  oxalate,  Mg02C202  •  2  H2O, 
(soluble  in  boiling  water  in  the  ratio  of  1  :  1300,  but  more 
soluble  in  presence  of  ammonium  salts)  is  likely  to  be  included 
in  the  precipitated  calcium  oxalate  and  in  proportion  to  the 
concentrations  of  both  the  calcium  salt  and  the  magnesium 
salt.  An  effective  procedure  suitable  in  presence  of  a  mag- 
nesium salt  *  depends  upon  bringing  about  the  precipitation  in 
the  dilute  and  boiling  solution  of  the  calcium  and  magnesium 
salts,  colored  with  methyl  orange  and  containing  a  considerable 
amount  of  ammonium  chloride,  by  adding  dropwise  at  intervals 
a  solution  of  oxalic  acid  (10  per  cent),  and  alternately  a  very 
dilute  ammonium  hydroxide  (1  per  cent)  to  neutralization. 
The  mixture  is  allowed  to  stand  for  an  hour  and  the  clear  liquid 
is  decanted  upon  an  ashless  filter.  In  case  the  magnesium 
present  is  very  largely  (more  than  ten-fold)  in  excess  of  the 
calcium,  the  precipitate  should  be  dissolved  in  hydrochloric 
acid,  reprecipitated  in  the  manner  described,  collected  upon 
the  filter  used  in  the  former  decantation,  and  washed  with  hot 
water.  The  paper  and  precipitate  are  ignited  in  a  platinum 
crucible,  cautiously  at  first  to  avoid  mechanical  loss  in  the 
evolution  of  the  gaseous  product,  afterward  at  the  full  heat  of 

*  Blasdale:    Analytical  Chemistry,   Treadwell-Hall,  3rd  ed.,  Vol.  II, 
p.  77. 


88  QUANTITATIVE  CHEMICAL  ANALYSIS 

the  burner,  and  finally  to  constant  weight  over  the  blast-lamp. 
The  residue  is  weighed  as  calcium  oxide,  CaO. 

Experimental  Process:  Determination  of  Calcium  Precipitated  as  Oxalate. 
—  Weigh  out  0.5  grm.  of  pure  dry  calcite  into  a  large  beaker,  add  a  little 
water  (20  cm.3),  cover  the  beaker  with  a  watch-glass,  and  add  concentrated 
hydrochloric  acid  gradually  until  the  evolution  of  carbon  dioxide  begins. 
Warm  gently  and  add  the  acid  until  solution  is  complete,  boil  to  expel 
carbon  dioxide,  neutralize  carefully  with  ammonium  hydroxide,  dilute  to 
a  volume  of  300  cm.3,  color  the  solution  with  methyl  orange,  and  heat 
to  boiling.  To  the  boiling  solution  add  dropwise  a  solution  of  1  grm.  of 
oxalic  acid  with  occasional  pauses  during  which  neutralization  is  slowly 
effected  by  means  of  1  per  cent  ammonium  hydroxide.  Let  the  mixture 
stand  an  hour  and  then  filter  upon  ashless  paper,  washing  with  hot  water. 
Ignite  carefully,  finally  at  the  full  heat  of  the  blast-lamp  and  weigh  as 
calcium  oxide,  CaO. 

The  Precipitation  of  Strontium  Oxalate.  —  The  appreciable 
solubility  of  strontium  oxalate  in  water  is  much  diminished  by 
the  addition  of  alcohol  to  the  solution.  When  the  alcohol 
added  amounts  to  a  fifth  of  that  of  the  water  solution  the 
precipitation  of  strontium  oxalate  by  an  excess  of  ammonium 
oxalate  is  practically  complete.  The  precipitate  thrown  down 
in  hot  solution  and  allowed  to  stand  over  night  may  be  filtered 
off  on  asbestos  in  the  perforated  crucible,  washed  with  20  per 
cent  alcohol,  ignited  gently  in  the  flame  of  the  Bunsen  burner, 
and  weighed  as  carbonate: 

SrCl2  +  (NH4)202C202  <=±  Sr02C202  +  2  NH4C1, 
SrO2C202  =  SrCO3  +  CO2. 

The  Precipitation  of  Barium  Oxalate.  —  The  addition  of 
alcohol  to  the  amount  of  a  fourth  of  the  water  solution  brings 
the  insolubility  of  barium  oxalate  within  the  range  of  solubilities 
permissible  in  fairly  good  analytical  operations.  An  excess  of 
ammonium  oxalate  added  to  the  alcoholic  solution  (25  per 
cent)  brings  about  a  practically  complete  precipitation  of  barium 
as  the  oxalate,  and  after  standing  over  night  the  precipitate 
may  be  collected  on  asbestos,  washed  with  25  per  cent  alcohol, 
ignited  gently,  and  weighed  as  carbonate: 

BaCl2  +  (NH4)202C202  <±  Ba02C202  +  2  NH4C1, 
Ba02C202  =  BaC03  +  C02. 

The  Precipitation  of  Mercurous  Oxalate.  —  Mercury  taken 
in  the  form  of  mercurous  nitrate  may  be  estimated  as  mer- 
curous  oxalate  precipitated  by  ammonium  oxalate: 

Hg2(N03)2  +  (NH4)202C202^Hg202C202  +  2NH4NO3. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  89 

It  is  necessary,  however,  to  control  the  acidity,  dilution,  and 
presence  of  mercuric  salts.  It  appears  that  5  cm.3  of  dilute 
nitric  acid,  sp.  gr.  1.15,  may  be  present  in  a  volume  of  100  cm.3 
and  that  5  cm.3  of  the  acid  will  prevent  precipitation  of  small 
amounts  of  mercuric  salt  (0.0100  grm.  to  0.0200  grm.,  calcu- 
lated as  mercury)  depending  upon  the  amount  of  ammonium 
oxalate  present  in  excess.  According  to  the  procedure  recom- 
mended,* mercurous  nitrate  dissolved  in  100  cm.3  of  water 
containing  2  to  5  per  cent  of  dilute  nitric  acid,  sp.  gr.  1.15, 
is  precipitated  by  the  addition  of  ammonium  oxalate  in  slight 
excess  with  stirring.  It  is  an  easy  matter  to  keep  the  excess 
of  the  precipitant  within  the  limits  of  1  cm.3  to  2  cm.3  of  the 
N/10  solution,  because  the  mercurous  oxalate,  when  properly 
stirred,  settles  rapidly.  The  precipitate  is  collected  on  asbestos 
in  a  perforated  crucible,  washed  two  or  three  times  with  cold 
water,  and  dried  to  constant  weight  over  sulphuric  acid^  at  the 
ordinary  temperature,  since  mercurous  oxalate  is  slowly  decom- 
posed at  temperatures  in  the  vicinity  of  100  degrees.  i 

The  Precipitation  of  Copper  Oxalate.  —  Copper  oxalate  is 
insoluble  in  water,  is  scarcely  attacked  by  moderate  amounts 
of  dilute  nitric  acid,  and  the  precipitation  of  copper  oxalate 
from  solutions  of  copper  salts  by  saturation  with  oxalic  acid 
is  very  nearly  complete,  provided  the  amount  of  the  copper 
salt  present  exceeds  the  amount  (equivalent  to  0.01  grm. 
of  copper  in  50  cm.3  of  liquid)  which  may  remain  in  super- 
saturated solution.  The  addition  of  a  large  proportion  of 
acetic  acid  to  the  solution  tends  to  prevent  supersaturation  as 
well  as  hydrolytic  decomposition,  and  small  amounts  as  well 
as  the  large  amounts  of  copper  may  be  successfully  precipitated 
as  the  oxalate  in  50  per  cent  acetic  acid.  The  precipitate  may 
be  collected  on  asbestos  in  the  perforated  crucible  and  con- 
verted by  careful  ignition  to  copper  oxide: 

CuS04  +  (NH4)202C262<^Cu02C202  +  (NH4)2S04. 


The  Precipitation  of  Carbonates 

In  precipitating  carbonates,  ammonium  carbonate  is  prefer- 
able to  sodium  carbonate  or  potassium  carbonate  whenever  its 
use  is  possible,  in  order  that  contamination  of  the  final  product 
by  inclusion  of  a  non-volatile  precipitant  may  be  obviated;  but 

*  Peters:  Methods  in  Analysis,  Gooch,  p.  195. 


90  QUANTITATIVE  CHEMICAL  ANALYSIS 

some  carbonates  are  not  quantitatively  precipitable  by  am- 
monium carbonate. 

The  Precipitation  of  Barium  Carbonate,  BaCO3;  Strontium 
Carbonate,  SrCO3;    and  Calcium  Carbonate,  CaCO3.  —  The 

carbonates  of  barium,  strontium,  and  calcium  may  be  pre- 
cipitated almost  completely  by  ammonium  carbonate  in  absence 
of  ammonium  salts  other  than  the  precipitant  and  the  salts 
produced  in  the  reaction.  Even  in  presence  of  ammonium 
salts  the  precipitation  may  be  made  practically  complete  by  a 
large  excess  of  ammonium  carbonate  in  a  50  per  cent  alcoholic 
solution: 


MC12  +  (NH4)2C03  <±  MC03  +  2  NH4C1. 

Experimental  Process:  Determination  of  Barium  as  Carbonate.  —  Dissolve 
0.5  grm.  of  barium  chloride,  BaC^  •  2  EkO  in  water  (50  cm.3),  add  am- 
monium carbonate  in  excess,  heat  the  liquid  to  granulate  the  precipitate, 
add  alcohol  until  it  equals  50  per  cent  of  the  solution,  settle  the  pre- 
cipitate, filter  on  asbestos  in  the  perforated  crucible,  wash,  ignite  to  low 
red  heat,  and  weigh  as  barium  carbonate,  BaCOs. 

The  Precipitation  of  Magnesium  Ammonium  Carbonate, 
MgCO3  •  (NH4)2CO3  •  4  H2O.  —  When  the  very  concentrated 
solution  of  the  sulphate,  nitrate,  or  chloride  of  magnesium  is 
treated  with  a  concentrated  solution  of  ammonium  carbonate 
a  voluminous  precipitate  forms  which  is  further  acted  upon  by 
an  excess  of  the  precipitant,  sometimes  dissolving  completely, 
and  is  converted  to  crystalline  magnesium  ammonium  carbon- 
ate, MgCO3  •  (NH4)2C03  •  4  H20.*  This  precipitate  is  notice- 
ably soluble  in  the  aqueous  solution  of  the  precipitant.  By 
adding  to  the  solution  an  equal  amount  of  alcohol  and  making 
the  precipitation  with  an  equal  amount  of  strongly  ammoni- 
acal  50  per  cent  alcohol  saturated  with  ammonium  carbonate 
the  precipitation  may  be  made  practically  complete.  f  In  prac- 
tice, the  solution  containing  salts  of  magnesium  (and  salts  of 
the  alkalies  besides)  is  brought  to  a  volume  of  about  50  cm.3, 
an  equal  amount  of  absolute  alcohol  is  added,  precipitation  is 

*  Schaffgotsch:  Ann.  Phys.,  104,  482  (1858). 

t  Gooch  and  Eddy:   Methods  in  Chemical  Analysis,  Gooch,  p.  154. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  91 

made  by  addition  of  50  cm.3  of  the  saturated  ammoniacal 
ammonium  carbonate  solution  containing  50  per  cent  alcohol, 
and,  after  stirring  for  five  minutes,  the  mixture  is  allowed  to 
strand  some  hours,  preferably  over  night.  If  an  alkali  salt  is 
present  in  amount  not  exceeding  0.1  grm.,  the  precipitate  may 
be  collected  on  asbestos  in  a  perforated  crucible,  washed  with  the 
precipitant,  dried,  ignited,  and  weighed  as  magnesium  oxide. 
When  the  amount  of  alkali  salt  originally  present  is  larger  the 
precipitate  may  be  freed  from  traces  of  the  alkali  salt  by  pour- 
ing off  the  supernatant  liquid  through  the  weighed  asbestos 
filter,  dissolving  the  precipitate,  and  precipitating  ammonium 
magnesium  carbonate  as  at  first.  This  second  precipitate, 
collected  upon  the  filter  originally  used,  leaves  upon  ignition 
practically  pure  magnesium  oxide: 

MgCl2  +  2  (NH4)2CO3  +  4  H2O  <=>  MgC03  •  (NH4)2C03  -  4  H2O 
+  2  NH4C1. 

Experimental  Process:  Determination  of  Magnesium.  —  Draw  from  a 
burette  50  cm.3  of  a  solution  of  magnesium  chloride  (about  N/10  in  re- 
spect to  the  anhydrous  salt)  standardized  as  the  sulphate,  according  to  the 
process  previously  described.  (Page  48.)  Add  an  equal  volume  of  ab- 
solute alcohol  (50  cm.3),  and  an  equal  volume  (50  cm.3)  of  saturated 
ammoniacal  solution  of  ammonium  carbonate  in  50  per  cent  alcohol.* 
Stir  for  five  minutes,  allow  the  mixture  to  stand  at  least  twenty  minutes, 
filter  on  asbestos,  wash  with  the  precipitant,  ignite,  and  weigh  the  residue 
as  magnesium  oxide,  MgO. 

The  Precipitation  of  Zinc  Carbonate  (Basic)  and  Cadmium 
Carbonate.  —  Zinc  carbonate  and  cadmium  carbonate  are  not 
completely  precipitated  by  ammonium  carbonate  nor  by  sodium 
or  potassium  carbonate  in  presence  of  ammonium  salts.  By 
treating  with  sodium  or  potassium  carbonate  the  solution  con- 
taining ammonium  salts  and  boiling,  the  ammonium  salts  may 
be  converted  into  sodium  salts  with  the  expulsion  of  ammonia. 
Upon  dissolving  in  hydrochloric  acid  the  precipitated  carbon- 
ate thus  formed,  which  may  include  considerable  amounts  of 
alkali  salt,  and  carefully  reprecipitating  with  the  fixed  alkali 

*  Best  made  by  adding  to  a  strong  aqueous  solution  of  ammonium  car- 
bonate, containing  ammonium  hydroxide,  an  equal  volume  of  absolute 
alcohol,  and  decanting  the  liquid  from  the  precipitate  of  ammonium  car- 
bonate thus  formed. 


92  QUANTITATIVE  CHEMICAL  ANALYSIS 

carbonate  the  precipitate  may  be  recovered  in  purer  condition. 
The  precipitant  is  added  to  the  acid  solution  until  a  turbidity 
is  produced,  the  solution  is  heated  to  precipitate  the  greater 
part  of  the  zinc  or  cadmium  carbonate  in  granular  form,  then 
a  few  drops  of  phenolphthalein  are  added,  and  enough  of  the 
precipitant  to  bring  out  a  pink  color  distinctly.  The  solution 
is  filtered  hot  on  asbestos  in  the  perforated  crucible,  washed 
with  hot  water,  ignited,  and  weighed  as  oxide.  In  this  process 
some  of  the  precipitate  falls  as  a  basic  salt  other  than  the 
carbonate  and  unless  this  basic  salt  is  completely  decomposed 
in  the  ignition  the  results  of  the  process  will  be  falsified.  For 
this  reason  the  process  is  at  its  best  when  applied  to  the  analysis 
of  nitrates  or  chlorides: 


3  Zn(N03)2  +  3  Na2C03  +  H20  +±  Zn3(OH)2(C03)2 
+  CO3  +  6  NaNO3. 

The  Precipitation  of  Bismuth  Basic  Carbonate.  —  Bismuth 
is  precipitated  from  the  solution  of  the  nitrate  by  addition  of 
ammonium  carbonate  in  slight  excess,  the  composition  of  the 
precipitate  being  variable  with  the  conditions  of  temperature, 
dilution,  and  excess  of  the  precipitant.  Ordinarily  the  pre- 
cipitate may  be  regarded  as  a  basic  carbonate  (hydroxy-oxy- 
carbonate),  approximating  the  constitution  represented  by  the 
symbol  Bi4(OH)203(C03)2;  but  with  the  basic  carbonate  is 
ordinarily  included  a  basic  bismuth  salt  derived  from  the 
soluble  bismuth  salt  acted  upon.  It  is  for  this  reason  that 
the  bismuth  salt  taken  for  analysis  by  this  method  should 
be  the  nitrate  which,  upon  ignition,  is  converted  to  oxide  as 
is  the  basic  carbonate.  The  precipitate  is  collected  upon  as- 
bestos in  the  perforated  crucible, 

12  BiCl3  +  18  (NH4)2CO3  +  3  H20 

<±  2  Bi2(CO3)3  •  3  Bi2O3  -  2  Bi03H3  +  12  C02  +  36  NH4C1, 

washed  with  hot  water,  ignited,   and  weighed  as  the  oxide, 
Bi203. 

The  Precipitation  of  Hydroxides 

Of  hydroxides  which  are  available  for  determinative  purposes, 
the  hydroxides  of  silicon  and  tungsten  may  be  taken  as  types 
of  acidic  hydroxides,  and  the  hydroxides  of  aluminium,  chro- 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  93 

mium,  and  iron  may  be  taken  as  representative  basic  hydroxides. 
The  acidic  hydroxides  may  be  formed  by  the  action  of  suitable 
acids  upon  susceptible  salts;  the  basic  hydroxides,  by  the 
action  of  soluble  hydroxides  or  by  the  hydrolytic  effect  of 
water  upon  susceptible  salts. 

Acidic  Hydroxides 

The  Precipitation  of  Silicic  Acid  from  Silicates.  —  When  a 
silicate  in  solution  is  treated  with  acid  the  silicic  acid  is  found 
as  hydrosol  or  as  hydrogel,  the  proportions  varying  according 
to  the  condition  of  action. 

Na2Si04  +  4  HC1  =  H4Si04  +  4  NaCl. 

Upon  evaporating  the  liquid  the  hydrosol  is  converted  to  the 
hydrogel  and  the  hydrogel,  losing  water,  becomes  more  in- 
soluble. Upon  desiccating  the  residue,  further,  best  at  about 
110  degrees,  the  silica  becomes  still  more  insoluble,  but  by  no 
means  absolutely  so;  for  upon  filtering  and  evaporating  the 
acidified  filtrate  appreciable  amounts  (some  milligrams),  de- 
pending upon  the  temperature  and  thoroughness  6f  the  desic- 
cation, may  be  recovered  in  a  second  filtration,  leaving  still 
smaller  amounts  which  may  be  separated  from  solution  by 
another  evaporation.  Repeated  moistenings  and  evaporations 
before  the  first  filtration  are  not  (in  the  presence  of  soluble 
salts)  as  effective  in  separating  the  silica  as  are  repeated  evapo- 
rations with  interpolated  filt rations.  A  final  ignition  at  red 
heat  leaves  the  silicon  dioxide  (silica)  in  anhydrous  form: 
H4Si04  =  SiO2  +  2  H20. 

Silica,  like  colloidal  precipitates  in  general,  is  known  to  take 
up  foreign  material  which  is  not  wholly  removed  when  the 
colloidal  form  is  modified  by  desiccation,  nor  by  digestion  with 
acid.  Upon  treating  such  a  precipitate  with  hydrofluoric  acid 
and  sulphuric  acid  silicon  fluoride  volatilizes  and  a  sulphate 
residue  is  left.  If  the  contaminating  substance  is  an  alkali 
salt,  like  the  sodium  chloride  resulting  from  the  action  of  acid 
upon  the  product  of  fusion  of  a  silicate  with  sodium  carbonate, 
the  residue  will  remain  as  sulphate  even  at  high  temperatures; 


94  QUANTITATIVE  CHEMICAL  ANALYSIS 

but  certain  other  contaminating  sulphates  may  on  strong 
ignition  leave  oxides.  It  is  plain,  therefore,  that  in  determining 
silica  separated  from  a  silicate  by  the  action  of  acid  the  amount 
of  included  impurity  should  always  be  determined  by  weighing 
the  residue  of  the  treatment  with  hydrofluoric  acid  and  sul- 
phuric acid;  and,  in  order  that  the  residue  of  this  treatment 
may  be  as  nearly  as  possible  identical  with  the  contaminating 
material  in  the  silica  when  it  is  weighed,  it  is  best  to  treat  the 
silica,  before  its  final  ignition  for  weighing,  with  a  drop  or  two 
of  strong  sulphuric  acid  and  to  subject  the  products  of  such 
treatment  to  a  temperature  similar  to  that  employed  in  the 
ignition  of  the  final  residue  left  after  the  treatment  with  hydro- 
fluoric acid  and  sulphuric  acid. 

Experimental  Process:  Determination  of  Silica.  —  Weigh  out  in  a  plati- 
num crucible  about  0.5  grm.  of  precipitated  silica,  the  purest  obtain- 
able.* Ignite  it  for  half  an  hour  over  the  full  heat  of  a  Bunsen  burner, 
cool,  and  weigh.  Take  the  difference  in  weights  as  the  weight  of  the 
anhydrous  silica.  Now  add  2  grm.  of  anhydrous  sodium  carbonate  and 
mix  it  thoroughly  with  the  silica  by  means  of  a  stiff  platinum  wire,  brush- 
ing off  into  the  crucible,  by  means  of  a  camel's  hair  brush,  any  particles  of 
material  adhering  to  the  wire.  Cover  the  crucible,  heat  the  mixture 
gradually  to  the  highest  temperature  of  the  Bunsen  burner  and  then,  if 
necessary  to  secure  complete  fusion,  .over  the  blast  lamp.  As  soon  as  the 
mass  is  in  quiet  fusion,  evolving  no  gas  bubbles,  so  that  it  may  be  assumed 
that  the  reaction  between  the  silica  and  the  sodium  carbonate  is  complete, 
take  the  crucible  in  tongs  applied  to  the  upper  edge,  and  give  it  a  gyratory 
motion  by  which  the  liquid  melt  is  spread  over  the  wall  of  the  crucible. 
Upon  cooling,  the  solid  material  may  be  loosened  from  the  crucible  by 
gentle  tapping.  Put  the  solid  material,  thus  loosened,  in  a  rather  tall 
beaker,  add  water  and,  gradually,  hydrochloric  acid  (sp.  gr.  1.1)  to  dis- 
solve the  excess  of  the  carbonate  and  to  break  up  the  silicate,  and  treat 
the  material  which  adheres  to  the  crucible  similarly.  In  case  the  greater 
part  of  the  material  adheres  obstinately,  place  the  crucible  in  the  beaker 
and  treat  it  with  the  adherent  material  in  the  manner  described.  Place 
the  beaker  on  a  water-bath  and  when  the  disintegration  of  the  melt  is 
complete  (aided  if  necessary  by  pressure  with  the  rounded  end  of  a  glass 
rod)  transfer  the  contents  to  a  dish  of  porcelain  or  (better)  platinum, 
evaporate  the  liquid,  and  dry  the  residue  at  110  degrees.  Drench  the 

*  The  purity  of  the  silica  may  be  tested  by  the  treatment  (with  hydro- 
fluoric acid  and  sulphuric  acid)  described  below  for  the  determination  of 
the  impurity  in  the  silica  obtained  in  the  analytical  process. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


95 


residue  with  concentrated  hydrochloric  acid,  add  an  equal  amount  of 
water,  cover  the  dish,  and  digest  for  half  an  hour,  with  occasional  stirring. 
Add  more  water,  stir  (and,  if  necessary,  grind  with  a  pestle  to  break  up 
lumps),  settle  the  precipitate,  pour  off  the  supernatant  liquid  upon  an 
ashless  filter,  wash  again  by  decantation,  transfer  the  precipitate  to  the 
filter,  wash,  and  dry. 

Return  the  filtrate  to  the  dish,  evaporate  to  dryness,  and  treat  the 
residue,  hi  the  manner  described  above,  for  the  recovery  of  the  silica 
(probably  several  milligrams)  which  did  not  separate  in  the  first  evapora- 
tion. Even  after  the  second  treatment,  traces  of  silica  escape  precipi- 
tation; but,  if  the  desiccation  at  110  degrees  has  been  thorough,  the 
amount  lost  should  not  exceed  about  0.0005  grm.,  in  the  average.  Put 
the  filter  containing  the  silica  in  a  weighed  crucible  recovered  in  the  second 
treatment  with  that  containing  the  main  portion  of  silica  and  ignite  cau- 
tiously until  the  paper  is  destroyed,  and  then  at  the  full  heat  of  the  Bunsen 
burner  for  half  an  hour.  Moisten  the  residue  with  a  drop  or  two  of  sul- 
phuric acid,  heat  cautiously,  ignite  to  low  redness,  cool,  and  weigh.  Then 
ignite  again  to  constant  weight.  The  difference  between  this  weight  and 
that  of  the  crucible  is  the  weight  of  the  silica  plus  some  sodium  sulphate. 

Moisten  the  silica  in  the  crucible  with  water,  add  10  cm.3  of  hydrofluoric 
acid,  which  distils  without  residue,  and  a  drop  or  two  of  concentrated 
sulphuric  acid,  and  evaporate  the  hydrofluoric  acid  (on  the  water-bath 
under  a  good  draft-hood)  repeating  the  operation  if  any  solid  residue  is  left. 
Continue  the  evaporation  to  remove  the  sulphuric  acid,  by  placing  the 
crucible  upon  a  triangle  suspended  within  a  large  porcelain  crucible,  or 
cone  of  iron  (the  radiator),  heating  the  latter  with  a  Bunsen  burner. 
Finally  ignite  the  crucible  and  residue  to  low  redness,  cool,  ignite,  and 
weigh. 

Deduct  the  weight  of  the  crucible  plus  that  of  the  residual  sodium 
sulphate,  which  is  inevitably  present,  from  the  previously  found  weight  of 
the  crucible,  silica,  and  residue  to  discover  the  weight  of  the  silica.  Com- 
pare the  weights  of  the  silica  found  with  that  of  the  silica  taken,  stating 
the  actual  results  and  the  percentage  result. 


Hydrous  silica 
taken. 

Anhydrous 
silica  taken. 

Silica  recovered 
with  impurity. 

Silica 
corrected. 

Error  in  recovery 
of  anhydrous 
silica. 

Grm. 

Grm. 

Grm. 

Grm. 

Grm. 

The  Determination  of  Tungsten   Trioxide  in   Tung  states.  — 
When  the  aqueous  solution  of  an  alkali  tungstate  is  treated 


96  QUANTITATIVE  CHEMICAL  ANALYSIS 

with  an  excess  of  hydrochloric  acid  and  submitted  to  evapo- 
ration the  tungstic  acid  is  very  largely  separated.  After  dry- 
ing at  120  degrees,  moistening  with  hydrochloric  acid,  and 
boiling  out  with  water  containing  hydrochloric  acid  (6  per  cent) 
or  ammonium  nitrate  (10  per  cent)  as  coagulants,  the  precipi- 
tate of  yellow  tungstic  acid  may  be  filtered  off,  ignited,  and 
weighed  as  tungsten  trioxide,  WO3.  But  as  in  the  determina- 
tion of  silica  by  similar  procedure  not  all  the  tungstic  acid  is 
recovered  in  one  operation.  Several  repetitions  of  the  process 
of  desiccation,  extraction,  and  filtration  are  essential  for  the 
complete  recovery  of  the  tungstic  acid: 

Na2W04  +  2  HC1  =  H2WO4  +  2  NaCl. 
H2W04  =  W03  +  H20. 

Basic  Hydroxides 

The  Precipitation  of  Aluminium  Hydroxide.  —  Aluminium 
hydroxide  is  easily  attacked  by  an  excess  of  the  fixed  alkali 
hydroxides,  and  these  reagents  are,  therefore,  not  available  for 
its  complete  precipitation.  Ammonium  hydroxide  in  slight 
excess  and  in  presence  of  an  ammonium  salt  (to  act  as  the 
coagulant  of  the  aluminium  hydroxide  hydrosol)  brings  about 
an  approximately  complete  precipitation  of  the  hydroxide  from 
solution  of  the  halogen  salts,  the  nitrate,  or  the  sulphate,  and 
precipitation  is  made  practically  complete  by  boiling  the 
solution. 

A1C13  +  3  NH4OH  <=±  A103H3  +  3  NH4C1. 

The  reaction  is  reversible  and  care  must  be  taken  to  make  sure 
that  the  solution  shall  always  retain  enough  of  the  volatile 
ammonia  to  render  it  distinctly  alkaline.  The  hydroxide  pre- 
cipitated in  this  manner  is,  'however,  not  entirely  pure,  but  re- 
tains some  of  the  acidic  ion  of  the  original  aluminium  salt  now 
combined  in  a  basic  salt,  which,  if  it  is  chloride,  nitrate,  or 
sulphate,  undergoes  dissociation  on  heating  and  leaves  the 
oxide,  as  does  the  hydroxide.  The  presence  of  certain  organic 
salts  —  e.g.,  tartrates  and  citrates  —  prevents  precipitation. 

The  precipitate  may  be  obtained  in  more  easily  filterable 
form  if  the  precipitation  is  first  made  by  a  hydrolytic  method, 
and  only  completed  by  the  direct  action  of  ammonium  hydrox- 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  97 

ide.  Thus,  the  precipitation  may  be  mainly  effected  by  treat- 
ing the  solution  of  the  aluminium  salt  with  sodium  thiosulphate 
(Chancel)  and  boiling  out  sulphur  dioxide: 

A12(S04)3  +  3  Na2S2O3  +  3  H2O  =  2  A103H3 

+  3  Na2SO4  +  3  S02  +  3  S. 

But  it  is  to  be  noted  that  in  this  operation  some  of  the  sulphur 
dioxide  may  become  oxidized  by  the  action  of  air  and  form 
sulphuric  acid, 

02  +  H20  =  2H2S04, 


so  that  the  precipitation  must  be  completed  by  addition  of  am- 
monium hydroxide  in  slight  excess  to  the  boiling  solution. 

Or  the  greater  part  of  the  precipitation  may  be  similarly 
brought  about  by  sodium  nitrite  (Wynkoop:  Schirm)  and  the 
neutralization  is  then  completed  by  ammonium  hydroxide: 

A12(SO4)3  +  6  NH4N02  +  3  H2O  =  2  A1203H3 

+  3  (NH4)  2  SO4  +  6  HNO2, 
3  HNO2  =  HN03  +  2  NO  +  H20, 
HN03  +  NH4OH  =  NH4N03  +  H20. 

Or  the  aluminium  salt,  which  of  itself  undergoes  incipient 
hydrolysis  in  water,  may  be  treated  with  an  iodide-iodate 
mixture  (which  at  once  reacts  with  the  acid  set  free),  liberat- 
ing iodine  and  the  iodine  taken  up  by  sodium  thiosulphate 
(Stock).  In  this  process  no  free  acid  is  formed  and  the  final 
addition  of  ammonium  hydroxide  is  not  necessary: 

2  AlCls  +  6  HOH  <=±  2  A103H3  +  6  HC1, 
KIO3  +  5  KI  +  6  HC1  =  6  KC1  +  3  H2O  +  3  12, 

I2  +  Na2S203  =  2  Nal  + 


Experimental  Processes:  Determination  of  Aluminium  Precipitated  as 
Hydroxide  and  Weighed  as  Oxide.  —  Dissolve  in  hot  water  about  1  grm. 
of  alum,  NH4A1(SO4)2  •  12  H2O  or  KA1(SO4)2  •  12  H2O,  weighed  exactly; 
dilute  the  solution  to  a  volume  of  about  300  cm.3;  add  concentrated 
nitric  acid  (about  5  cm.3),  to  make  a  considerable  amount  of  ammonium 
salt  in  the  subsequent  neutralization  and  thus  favor  the  formation  of  the 
aluminium  hydroxide  hydrogel  rather  than  the  hydrosol;  heat  the  solu- 
tion to  boiling  and  add  in  slight  excess  ammonium  hydroxide  freshly  made 
by  charging  distilled  water  with  ammonia  gas,  in  order  that  it  may  be 
free  from  contaminating  materials  taken  up  by  long  standing  in  glass; 
and  set  the  mixture  aside  to  settle  the  precipitate.  Decant  the  clear 
supernatant  liquid  upon  a  paper  filter  protected  at  the  point  by  a  per- 
forated cone  of  platinum  or  of  parchment  paper;  dilute  the  remaining 


98  QUANTITATIVE  CHEMICAL  ANALYSIS 

liquid  to  the  original  volume  with  hot  water  containing  ammonium  nitrate 
(1  per  cent)  and  a  very  little  ammonium  hydroxide;  transfer  the  liquid  to 
the  filter;  and  wash  with  slightly  ammoniacal  hot  water  containing 
ammonium  nitrate  (1  per  cent).  At  the  end  of  the  washing,  apply  the 
suction  of  the  filter  pump  to  compact  the  precipitate;  dry,  ignite  cau- 
tiously at  first,  and  later  at  the  full  heat  of  the  Bunsen  burner,  until  the 
carbon  of  the  paper  has  been  entirely  consumed  and  any  basic  aluminium 
sulphate  of  the  precipitate  has  been  converted  to  oxide,  and  weigh  the 
residue  as  A12O3. 

The  Precipitation  of  Chromic  Hydroxide.  —  Chromium  in 
chromic  chloride,  nitrate  or  sulphate  is  precipitated  by  the 
methods  described  above  for  the  precipitation  of  aluminium 
hydroxide,  and  the  reactions  involved  are  precisely  similar. 
From  a  soluble  chromate  the  chromic  salt  may  be  formed  by 
the  action  of  hydrochloric  acid  and  alcohol  in  the  boiling  solu- 
tion (with  formation  of  aldehyde), 

K2Cr2O7  +  14  HC1  +  3  C2H5OH 

=  2  KC1  +  2  CrCl3  +  3  C12  +  7  H20  -f  3  C2H5OH 
=  2  KC1  +  2  CrCl3  +  7  H2O  +  3  C2H4O  +  6  HC1, 

or  by  sulphur  dioxide, 

K2Cr207  +  3  S02  +  H2S04  =  K2S04  +  Cr2(S04)3  +  H20, 

and  the  chromic  chloride  or  sulphate  may  then  be  precipitated 
as  hydroxide,  ignited,  and  converted  to  oxide  for  weighing. 

Experimental  Process:  Determination  of  Chromium  Precipitated  as  Chromic 
Hydroxide  and  Weighed  as  Chromic  Oxide.  —  Dissolve  in  a  porcelain 
dish  about  0.5  grm.  of  potassium  dichromate  (weighed  exactly)  in  10  cm.3 
of  hot  water,  add  15  cm.3  of  alcohol,  and  5  cm.3  of  concentrated  hydro- 
chloric acid.  Cover  with  a  water-glass  (to  prevent  loss  by  spattering)  and 
boil  until  the  clear  green  color  of  the  liquid  indicates  that  the  reduction 
has  been  completed.  Dilute  the  solution  to  a  suitable  volume  (250  cm.3) 
with  hot  water,  add  ammonium  hydroxide  in  very  slight  excess  (or  make 
the  neutralization  nearly  complete  with  ammonium  hydroxide  and  com- 
plete the  precipitation  by  means  of  freshly  made  ammonium  sulphide 
added  in  slight  excess) ;  heat  the  mixture  to  boiling  until  the  liquid  above 
the  precipitate,  allowed  to  settle,  is  colorless.  Filter  on  ashless  paper, 
wash  carefully  with  hot  water.  Dry  precipitate  and  paper,  separate  the 
precipitate  as  well  as  possible  from  the  paper,  put  the  precipitate  in  a  plati- 
num crucible,  burn  the  paper  held  in  a  spiral  of  platinum  wire  over  the 
crucible,  and  add  the  ash  to  the  contents  of  the  crucible.  Ignite  gently  at 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  99 

first,  then  at  red  heat  in  a  draft  of  air  to  constant  weight.  Weigh  the 
residue  as  Cr2O3. 

The  Precipitation  of  Ferric  Hydroxide.  —  Iron  in  the  form 
of  ferric  chloride,  nitrate,  or  sulphate  is  precipitable  by  am- 
monium hydroxide  in  presence  of  ammonium  salts  and  in  this 
case  the  presence  of  a  considerable  excess  of  the  precipitant 
does  not  materially  increase  the  solubility  of  the  precipitate: 

FeClg  +  3  NH4OH  +±  Fe03H3  +  3  NH4C1. 

The  precipitate  retains  some  of  the  acidic  ion  in  the  form  of  a 
basic  salt,  but  this,  if  it  is  a  chloride,  nitrate,  or  sulphate,  is 
dissociated  on  ignition.  Like  aluminic  hydroxide,  ferric  hydrox- 
ide (or  a  basic  salt)  is  also  precipitable  by  the  iodide-iodate 
mixture  or  by  ammonium  nitrite;  but  sodium  thiosulphate  re- 
duces the  ferric  salt  to  the  ferrous  condition  and  fails  to  bring 
about  precipitation.  In  fact,  aluminium  and  iron  may  be 
separated  approximately  by  the  hydrolytic  action  of  sodium 
thiosulphate  upon  the  mixed  salts.  Iron  in  a  ferrous  salt  may 
be  determined  by  oxidation  with  a  suitable  oxidizer  (chlorine, 
bromine,  or  nitric  acid,  or  the  iodide-iodate  mixture)  and 
treatment  of  the  ferric  salt  as  described. 

Experimental  Process:  Determination  of  Iron  Precipitated  as  Ferric 
Hydroxide  and  Weighed  as  Ferric  Oxide.  —  Weigh  out  exactly  about  1  grm. 
of  clean  crystals  of  ferrous  ammonium  sulphate,  Fe(NH4)2(SO4)2  •  6H2O; 
dissolve  in  water  (100  cm.3),  add  hydrochloric  acid  (5  cm.3),  and  nitric  acid 
(2  cm.3).  Boil  the  solution  (with  addition  of  more  nitric  acid,  if  necessary) 
until  the  precipitate  formed  on  addition  of  a  drop  or  two  of  ammonia  is 
reddish  —  not  black.  Dilute  the  solution  to  a  convenient  volume  (300 
cm.3),  add  ammonium  hydroxide  in  slight  excess,  boil,  filter  on  paper,  wash 
with  care,  and  ignite. 

The  Basic  Acetate  Process.  —  The  acetates  of  iron  and 
aluminium  are  hydrolyzed  when  heated  in  dilute  solution  and 
if  a  sufficient  amount  of  a  coagulant  be  present  the  hydroxides 
or  basic  salts  may  be  precipitated  in  the  hydrogel  form.  On 
the  other  hand,  acetates  of  nickel,  cobalt,  zinc,  and  manga- 
nese -are  not  so  easily  hydrolyzed.  The  basic  acetate  precipi- 
tation may  therefore  be  applied  to  the  separation  of  iron  and 
aluminium  from  nickel,  cobalt,  zinc,  and  manganese  as  well  as 
calcium,  barium,  strontium,  magnesium,  and  the  akali  elements, 


100  QUANTITATIVE  CHEMICAL  ANALYSIS 

although  for  aluminium  unaccompanied  by  iron  the  separation 
is  less  satisfactory  than  for  iron  or  for  iron  with  aluminium. 

After  neutralizing  the  solution  with  great  care  by  means  of 
sodium  carbonate,  or  ammonium  carbonate,  added  dropwise 
in  dilute  solution  until,  after  vigorous  stirring,  a  faint  per- 
manent turbidity  remains,  hydrochloric  acid  is  added  dropwise, 
with  long  intervals  between  the  additions  and  stirring,  until 
the  turbidity  (which,  if  the  solution  is  dark  from  much  iron, 
may  be  better  seen  against  a  white  background)  is  just  dis- 
solved. To  the  clear  solution  is  then  added  2  grm.  to  3  grm.  of 
sodium  acetate,  or  ammonium  acetate  (made  neutral),  in  solu- 
tion, and  a  large  volume  (400  cm.3  to  700  cm.3)  of  boiling  water. 
The  liquid  is  boiled  for  some  minutes,  and  then  filtered  hot 
without  suction  upon  a  filter  protected  for  the  subsequent  ap- 
plication of  suction,  and  washed  with  hot  water  containing  a 
little  sodium  acetate  or  ammonium  acetate.  After  the  wash- 
ing is  complete  suction  is  applied  to  compact  the  precipitate. 

The  precipitate  obtained  in  the  manner  described  may  still 
retain  traces  of  nickel,  cobalt,  zinc,  and  manganese,  and  traces 
of  aluminium  and  iron  will  have  been  carried  into  the  filtrate. 
The  precipitate  is,  therefore,  dissolved  in  hot  hydrochloric  acid 
(1:2)  and  reprecipitated  by  ammonium  hydroxide  to  recover 
the  iron  hydroxide  and  aluminium  hydroxide  in  purer  con- 
dition; and  from  the  combined  filtrates  the  dissolved  traces  of 
these  two  hydroxides  are  recovered  by  concentration  to  the  low- 
est convenient  limit,  treatment  with  a  bare  excess  of  ammo- 
nium hydroxide,  and  filtration.* 

The  Precipitation  of  Sulphides 

Some  insoluble  sulphides  are  precipitable  by  hydrogen  sul- 
phide from  water  solution  in  the  presence  of  free  acid,  and  some 
only  in  alkaline  solution. f  Precipitated  sulphides,  however, 

*  It  should  be  noted  that  the  precipitates  found  in  the  presence  of 
phosphoric  acid  and  arsenic  acid  contain  aluminiac  and  ferric  salts  of  these 
acids. 

t  In  precipitating  the  latter  ammonium  sulphide  is  used  as  the  precipi- 
tant, except  in  certain  special  cases.  Such  precipitations  are  often  made 
to  effect  the  separation  of  the  precipitable  sulphides  from  salts  of  the 
alkali-earth  elements;  and  in  these  cases  it  is  essential  that  the  ammonium 
sulphide  should  be  freshly  made  and  free  from  ammonium  carbonate. 
The  ammonium  sulphide  may  be  prepared  free  from  carbonate  by  satu- 
rating with  hydrogen  sulphide  a  solution  of  ammonium  hydroxide  which 
has  been  made  by  absorbing  in  water,  freshly  boiled  and  cooled  in  a  cur- 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS 


101 


are  generally  oxidizable  in  the  air  at  moderately  elevated  tem- 
peratures, and  if  they  are  to  be  weighed  as  such,  must  be  pro- 
tected from  oxidation  in  the  preparation  for  weighing,  by 
taking  care  to  moderate  the  temperature  of  desiccation  (as  in 
drying  mercury  sulphide,  and  the  arsenic  sulphides),  by  heat- 
ing (Fig.  21)  in  a  current  of  pure  carbon  dioxide  (as  in  deter- 


FIG.  21. 


FIG.  22. 


mining  antimony  trisulphide,  see  p.  104),  or  by  ignition  in 
mixture  with  pure  sulphur  in  a  Rose  crucible  (Fig.  22)  kept 
filled  with  hydrogen  which  is  introduced  through  the  perforated 
cover  (as  in  determining  copper,  see  p.  105).  In  some  cases  it 
is  necessary  to  convert  the  precipitated  sulphide  to  some  other 
form  for  weighing. 

Sulphides  Precipitable  in  Acid  Solution 

The  Precipitation  of  Mercuric  Sulphide  by  Hydrogen 
Sulphide.  —  The  formation  of  insoluble  mercuric  sulphide  by 
the  action  of  hydrogen  sulphide  upon  the  mercuric  salt  in  solu- 
tion is  a  simple  process  well  adapted  to  the  determination  of 
mercury,  provided  the  solution  does  not  contain  nitric  acid  or 
other  oxidizer  which  will  set  free  sulphur  from  hydrogen  sul- 
phide. In  the  first  action  complex  double  compounds  of  light 
color  are  formed  which  are  converted  to  the  black  mercuric 

rent  of  purified  air,  the  gaseous  ammonia  evolved  from  a  solution  of  con- 
centrated ammonium  hydroxide  first  digested  with  lime  and  then  brought 
to  the  boiling  point. 


102  QUANTITATIVE  CHEMICAL  ANALYSIS 

sulphide  when  the  liquid  is  thoroughly  saturated  with  hydrogen 
sulphide. 

(m  +  n)HgC!2  +  wH2S  =  wHgS  -  nHgC!2  +  2  wHCl. 
mHgS  •  nHgCl2  +  nH2S  =  (m  +  w)HgS  +  2  nHCl. 

The  precipitate,  after  settling,  is  filtered  on  asbestos  in  the  per- 
forated crucible,  washed  with  cold  water,  dried  at  105  to  110 
degrees  and  weighed  as  mercuric  sulphide. 

Experimental  Process:  Determination  of  Mercury  in  Mercuric  Chloride. 
Dissolve  0.5  grm.  of  mercuric  chloride  in  a  little  hot  water,  dilute  to 
100  cm.3  with  cold  water,  and  saturate  the  cold  solution  with  hydrogen 
sulphide.  Collect  the  precipitate  upon  asbestos  in  the  perforated  crucible, 
dry  at  105  to  1 10  degrees  to  constant  weight,  and  weigh  as  mercuric  sul- 
phide, HgS. 

When  mercuric  nitrate  is  treated  with  hydrogen  sulphide  the 
sulphur  set  free  in  a  secondary  action  between  nitric  acid  and 
hydrogen  sulphide  may  contaminate  the  precipitated  sulphide 
and  vitiate  the  result;  and  this  effect  is  much  magnified  when 
the  solution  contains  a  considerable  excess  of  nitric  acid,  as  is 
often  the  case  in  practical  analysis.  In  such  cases  the  following 
more  elaborate  procedure  is  adopted. 

The  Precipitation  of  Mercuric  Sulphide  by  Decomposition 
of  the  Sulphosalt.  —  The  solution  of  the  mercuric  salt  acid  in 
reaction  is  treated  with  sodium  carbonate  nearly  to  neutraliza- 
tion, and  then  a  mixture  of  freshly  prepared  ammonium  sul- 
phide and  sodium  hydroxide  is  added,  with  warming  and 
stirring,  until  the  soluble  mercury  sulphosalt  is  formed.  Am- 
monium nitrate  is  added  and  the  solution  is  boiled  until  the  free 
ammonia  is  nearly  all  expelled: 

Hg(N03)2  +  4  (NH4)2S  +  4  NaOH 

=  HgS  +  Na2S  +  2  NaNO3  +  4  NH4OH 
=  Hg(SNa)2  +  2  NaNO3  +  4  NH3  +  4  H20, 

Hg(SNa)2  +  2  NH4N03  =  HgS  +  2  NaNO3  +  (NH4)2S. 

The  precipitated  mercuric  sulphide  is  settled,  collected  on  as- 
bestos in  the  perforated  crucible,  and  washed  with  hot  water; 
and  admixed  sulphur  is  extracted  by  pure  carbon  disulphide 
applied  after  washing  away  the  repellant  water  from  the  pre- 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS          103 

cipitate  by  means  of  alcohol.  A  final  washing  with  alcohol, 
followed  by  ether,  prepares  the  precipitate  for  drying  at  105  to 
110  degrees  in  the  air-bath. 

The  Precipitation  of  Arsenic  Trisulphide,  As2S3. —  When 
the  cold  solution  of  compounds  of  arsenic  in  the  trivalent  con- 
dition is  acidified  strongly  with  hydrochloric  acid,  hydrogen 
sulphide  precipitates  arsenic  trisulphide: 

2  AsCl3  +  3  H2S  <±  As2S3  +  6  HC1. 

The  excess  of  hydrogen  sulphide  is  largely  removed  from  the 
solution  by  a  current  of  carbon  dioxide  and  the  arsenic  tri- 
sulphide, collected  upon  asbestos  in  the  perforated  crucible,  is 
dried  to  constant  weight  at  105  degrees. 

The  Precipitation  of  Arsenic  Pentasulphide,  As2S5.  —  If  the 
solution  of  compounds  of  arsenic  in  quinquivalent  form  is  kept 
cool  (best  by  immersing  the  container  in  ice-water)  while  being 
treated  with  twice  its  volume  of  concentrated  hydrochloric  acid, 
the  arsenic  remains  in  solution  as  the  pentachloride.  By 
rapidly  saturating  such  a  solution  with  hydrogen  sulphide,  and 
allowing  it  to  stand  in  a  stoppered  flask  for  a  considerable 
period  (two  hours)  arsenic  pentasulphide  is  precipitated: 
2  AsCl3  +  5  H2S  <F±  As2S5  +  10  HC1. 

This  precipitate  may  be  collected  on  asbestos  in  the  perforated 
crucible,  washed  with  water  and  then  with  alcohol  (to  facili- 
tate the  removal  of  water),  dried  at  105  degrees,  and  weighed 
as  the  pentasulphide.  It  is  essential,  however,  that  the  acidi- 
fied solution  be  kept  thoroughly  cooled  until  the  pentasul- 
phide is  precipitated,  else  the  higher  chloride  will  dissociate 
partially  to  arsenic  trichloride  and  chlorine  with  the  result  that 
the  precipitate  will  consist  of  arsenic  trisulphide,  arsenic 
pentasulphide,  and  sulphur  in  indefinite  mixture,  while  some  of 
the  sulphur  may  be  lost  in  the  filtration  or  undergo  oxidation 
in  the  drying. 

The  Precipitation  of  Antimony  Trisulphide,  Sb2S3.  — 
Antimonous  Salts.  When  a  current  of  hydrogen  sulphide  acts 
upon  a  salt  of  trivalent  antimony  in  a  hot  solution,  sufficiently 
acidulated  with  hydrochloric  acid  (with  addition  of  tartaric 


104  QUANTITATIVE  CHEMICAL  ANALYSIS 

acid  or  without  it,  as  may  be  preferred),  to  prevent  the  precipi- 
tation of  an  antimony  oxychloride,  antimony  trisulphide  is 
precipitated  as  an  orange  precipitate  which  grows  redder  and 
denser  if  the  solution  is  kept  at  the  boiling  temperature  while 
the  current  of  hydrogen  sulphide  is  passed  for  a  considerable 
time: 

2  SbCl3  +  3  H2S  <±  Sb2S3  +  6  HC1. 

The  precipitate  may  be  settled,  collected  upon  asbestos  in  a 
perforated  crucible,  and  washed  with  a  hot  1  per  cent  (acid) 
solution  of  ammonium  acetate  containing  hydrogen  sulphide, 
followed  in  succession  by  alcohol,  carbon  disulphide,  alcohol, 
and  ether,  and  dried  at  110  degrees  for  weighing. 

If  the  boiling  solution  in  which  precipitation  takes  place  is 
made  up  to  contain  a  fourth  of  its  volume  of  concentrated 
hydrochloric  acid,  the  precipitate  will  become  very  dark,  dense 
and  crystalline  upon  prolonged  boiling  in  the  current  of  hydro- 
gen sulphide  and  agitation  (Vortmann  and  Metzel) .  Under  these 
conditions  the  precipitation  is  not  quite  complete  but  may  be 
made  so  by  diluting  the  solution  with  an  equal  volume  of  water 
while  continuing  the  operation  a  few  minutes  longer.  The  pre- 
cipitate may  be  collected  upon  asbestos,  and  washed  with  hot 
water  (without  showing  any  tendency  to  pass  into  the  hydrosol 
form)  followed  in  succession  by  alcohol,  carbon  disulphide, 
alcohol,  and  ether,  and  dried  at  110  degrees;  or  the  precipi- 
tated antimony  suphide  may,  after  the  washing  with  hot  water, 
be  heated  in  carbon  dioxide,  as  described  below. 

Antimonic  Salts.  —  By  a  procedure  similar  to  that  which  has 
been  described  as  applicable  to  the  precipitation  of  antimony 
trisulphide  from  antimonous  salts,  the  same  product,  mixed  with 
sulphur,  may  be  obtained  from  antimonic  salts: 

2  SbCl5  +  5  H20  <±  Sb2S3  +  10  HC1  =  Sb2S3  +  10  HC1  +S2. 

In  this  case  the  removal  of  the  excess  of  sulphur  may  be  best 
accomplished  by  volatilization.  The  precipitate  is  washed  with 
aqueous  ammonium  acetate,  or  with  water,  heated  in  the  filter 
crucible  enclosed  in  a  tube  or  other  suitable  apparatus  (see  p. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  105 

101),  in  an  air-free  atmosphere  of  carbon  dioxide  first  at  100  to 
130  degrees,  and  finally  at  270  to  280  degrees,  and  weighed  as 
antimony  trisulphide. 

The  Precipitation  of  Cupric  Sulphide  and  Weighing  as 
Cuprous  Sulphide.  —  In  order  that  precipitated  cupric  sul- 
phide may  settle  well  it  is  precipitated  from  a  solution  con- 
taining a  regulated  amount  of  free  acid.  The  solution  should 
contain  5  per  cent  by  volume  of  concentrated  acid,  preferably 
sulphuric  acid.  Hydrogen  sulphide  is  passed  into  the  hot  solu- 
tion and  the  current  of  gas  is  continued  while  the  solution 
cools.  The  precipitate  thus  formed, 

CuS04  +  H2S  +±  CuS  +  H2SO4, 

is  settled,  collected  upon  a  paper  filter  protected  by  a  platinum 
point  so  that  suction  may  be  later  applied  without  danger  to 
the  filter,  although  at  the  outset  the  filtration  is  best  made 
without  suction.  It  is  washed  with  water  charged  with  hydro- 
gen sulphide  and  containing  a  little  acetic  acid.  The  filter 
should  be  kept  full  to  the  very  last  during  the  filtration  and 
washing,  to  avoid  oxidation  of  the  sulphide  to  soluble  form 
(sulphate  and  thiosulphate)  which  by  reaction  with  the  hydrogen 
sulphide  will  again  form  cupric  sulphide  in  the  hydrosol  con- 
dition precipitable  by  coagulation  in  the  filtrate.  Finally,  suc- 
tion is  applied  to  drain  the  precipitate  which  is  dried  on  the 
paper  at  90  to  100  degrees.  The  precipitate  is  transferred  as 
completely  as  may  be  to  a  Rose  crucible  (p.  101) ;  the  paper 
is  burned  in  a  platinum  spiral,  and  the  ash,  first  collected  on  a 
crucible  cover  and  reignited  to  burn  all  carbon,  is  introduced; 
pure  sulphur  (recrystallized)  from  carbon  disulphide,  so  that 
it  will  leave  no  residue  when  volatilized,  is  added;  the  cover 
is  put  in  place  upon  the  crucible,  and  through  the  perforation 
hydrogen  is  introduced  from  a  generator  at  the  rate  of  about 
four  bubbles  per  second  in  the  wash  bottle;  the  crucible  is 
heated,  gently  at  first  and  then  so  that  the  bottom  shows  faint 
redness,  until  the  hydrogen  flame  shows  no  blue  color  of  sulphur 
and  emits  no  odor  of  sulphur  dioxide: 

2  CuS  =  Cu2S  +  S. 


106  QUANTITATIVE  CHEMICAL  ANALYSIS 

The  crucible  is  cooled  nearly  to  atmospheric  temperature  in  an 
increased  flow  of  hydrogen,  placed  in  a  desiccator  to  cool 
thoroughly,  and  weighed.  The  difference  between  the  weight 
of  the  crucible  and  that  of  the  crucible  with  the  residue  is  the 
weight  of  the  cuprous  sulphide,  Cu2S. 

The  Precipitation  of  Zinc  Sulphide.  —  Zinc  may  be  precipi- 
tated as  the  sulphide  from  ammoniacal  solutions  or  from  solu- 
tions of  regulated  acidity. 

A.  According  to  the  first  procedure  the  solution,  contained 
in  an  Erlenmeyer  flask,  is  treated  with  sodium  carbonate  until 
some  zinc  carbonate  is  precipitated.  The  precipitate  is  dis- 
solved by  a  very  little  ammonium  hydroxide  and  to  the  solution 
is  added  ammonium  acetate,  or,  better,  ammonium  thiocyanate, 
in  the  proportion  of  5  grm.  to  every  100  cm.3  (to  coagulate  the 
.sulphide  to  be  formed),  and  then  freshly  made  ammonium 
.sulphide  in  excess: 

ZnCl2  +  (NH4)2S  <=±  ZnS  +  2  NH4C1. 

The  mixture  is  diluted  with  boiled  water,  and  the  flask,  nearly 
full  and  stoppered,  is  set  aside  to  stand  some  hours  (twelve  to 
twenty-four).  The  clear  liquid  is  decanted  upon  a  filter;  the 
precipitate  is  diffused  through  a  5  per  cent  solution  of  am- 
monium acetate  or  thiocyanate  and  again  settled;  and  the 
liquid  is  poured  upon  the  same  filter  while  the  filtrate  is  allowed 
to  run  into  an  empty  beaker  so  that  in  case  of  necessity  it  may 
be  refiltered  apart  from  the  first  decant ation.  After  several 
repetitions  of  the  washing  by  decantation  the  precipitate  is 
transferred  to  the  filter,  carefully  washed  with  the  coagulating 
mixture  (keeping  the  filter  full)  and  finally  with  water  con- 
taining ammonium  sulphide.  The  dried  precipitate  and  the 
filter  may  be  ignited  apart,  and  the  combined  residues  mixed 
with  sulphur,  ignited  in  hydrogen  (see  p.  101),  and  weighed  as 
zinc  sulphide,  ZnS.  Or  the  precipitated  sulphide  may  be  dis- 
solved-from  the  filter  in  hot  dilute  hydrochloric  acid;  the 
solution  received  and  evaporated  to  dryness  in  a  weighed 
crucible;  the  residue  dissolved  in  a  little  water,  mixed  with 
pure  (precipitated)  mercuric  oxide  (Volhard) ;  the  mixture  dried 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS          107 

and  gently  ignited  to  volatilize  the  mercuric  chloride  formed 
and  the  mercury  derived  from  the  excess  of  mercuric  oxide;  and 
the  residue  weighed  as  zinc  oxide: 

ZnS  +  2  HC1  =  ZnCl2  +  H2S, 
ZnCl2  +    HgO  =  ZnO    +  HgCl2.       . 

B.  According  to  a  second  procedure  (applicable  to  the  sepa- 
ration of  zinc  from  alkali  and  alkali-earth  elements)  ammonium 
chloride  or  sulphate   (5  grm.   to   100  cm.3)   and  ammonium 
acetate  (1  grm.)  are  added  to  the  slightly  acid  solution  and 
hydrogen  sulphide  is  passed  into  the  solution  to   saturation. 
The  precipitate,  which  coagulates  much  better  in  presence  of 
the  acetic  acid  set  free  than  in  ammoniacal  solution,  is  allowed 
to  settle  completely  and  then  is  collected  upon  paper  and 
washed  with  a  two  per  cent  solution  of  acetic  acid  saturated 
with   hydrogen   sulphide;   and  the  paper  and   precipitate  are 
treated  as  described  in  the  preceding  paragraph: 

Zn02  (C2H30)2  •+  H2S  <=»  ZnS  +  2  HOC2H3O. 

C.  To  precipitate  zinc  as  sulphide  while  nickel  and  cobalt 
salts  remain  in  solution  requires  the  careful  adjustment  of  a 
notable  excess  of  acid.     Instead  of  depending  upon  regulation 
of  the  amount  of  free  acetic  acid  (as  may  also  be  done  more 
or  less  successfully)  it  is  better  to  rely  upon  a  limited  excess 
of  hydrochloric  acid  to  prevent  precipitation  of  the  nickel  and 
cobalt  sulphides  and  then  to  take  advantage  of  the  coagulating 
action  of  the  neutral  salt  of  a  strong  acid  to  precipitate  the 
colloidal  zinc  sulphide. 

To  the  acid  solution  from  which  zinc  sulphide  is  to  be  precipi- 
tated, sodium  carbonate  is  added  until  a  slight  permanent 
precipitate  is  formed.  Hydrochloric  acid  (N/10)  is  added,  with 
stirring,  until  the  precipitate  is  just  dissolved  and  then  in  ex- 
cess amounting  to  10  cm.3  of  the  N/10  acid  (0.0365  grm.)  in 
every  100  cm.3.  Next  is  added  an  electrolyte  (2  grm.  to  5  grm. 
in  100  cm.3)  to  serve  as  a  coagulant  of  the  sulphide  to  be  formed, 
best  ammonium  thiocyanate  (Zimmermann) ,  chloride,  or  sul- 
phate (Kramers);  the  solution  is  heated  to  50  or  60  degrees; 


108  QUANTITATIVE  CHEMICAL  ANALYSIS 

saturated  with  hydrogen  sulphide,  and  allowed  to  stand  until 
the  pure  white  precipitate  has  settled: 
ZnCl2  +  NiCl2  +  CoCl2  +  n(NH4Cl)  +  H2S 

<=±  ZnS  +  NiCl2  +  CoCl2  +  nNH4Cl  +  2  HC1. 

The  precipitate  is  collected  on  paper,  washed  with  a  solution 
of  the  coagulant  (2  per  cent),  dried,  ignited  until  the  paper  is 
ashed,  reignited  with  sulphur  in  hydrogen  and  weighed  as 
sulphide;  or  it  may  be  converted  to  oxide  and  weighed  as  in  A. 
Precipitation  of  Nickel  Sulphide  and  Cobalt  Sulphide  in 
Acetic  Acid  Solution.  —  Nickel  and  cobalt  may  be  precipitated 
nearly  completely  from  a  solution  which  is  acid  only  with 
acetic  acid  at  not  too  great  concentration: 

Ni02(C2H30)2  +  H2S  <=>  NiS  +  2  HOC2H30, 
Co02(C2H30)  +  H2S  <=>  CoS  +  2  HOC2H30. 

The  process  is  adapted  to  the  separation  of  these  elements 
from  manganese  as  well  as  from  the  alkali  and  alkali-earth 
elements.  The  solution  of  small  volume  containing  chlorides 
or  sulphates  of  these  elements  is  treated  with  an  excess  of 
sodium  carbonate,  then  made  strongly  acid  with  acetic  acid, 
diluted  to  a  volume  of  100  to  200  cm.3,  treated  with  ammonium 
acetate  (amounting  to  5  per  cent  of  the  solution),  heated  to 
70  or  80  degrees,  and  charged  to  saturation  with  hydrogen 
sulphide.  The  precipitate  of  nickel  sulphide  and  cobalt  sul- 
phide is  filtered  off  and  washed  with  hot  water.  In  the  filtrate 
will  be  found  the  manganese,  the  alkali  and  alkali-earth  ele- 
ments, with  small  amounts  of  nickel  and  cobalt  which  may  be 
recovered  by  concentrating  the  solution,  neutralizing  with 
ammonium  hydroxide  containing  a  little  fresh  ammonium 
sulphide,  making  slightly  acid  with  acetic  acid,  warming,  and 
again  filtering.  The  recovered  traces  and  the  main  precipitate 
of  nickel  sulphide  and  cobalt  sulphide  are  dried  and  ignited, 
with  the  filters,  and  the  residue  of  mixed  oxides  and  sulphides 
is  dissolved  in  aqua  regia.  Repeated  evaporations  with  hydro- 
chloric acid  leave  a  residue  of  chlorides  which  may  be  treated 
by  appropriate  methods  (see  p.  112)  for  the  separation  and 
determination  of  nickel  and  cobalt. 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS          109 

The  Precipitation  of  Stannous  Sulphide  and  Stannic  Sul- 
phide, with  Weighing  as  Stannic  Oxide.  — From  stannous 
chloride  in  solution  hydrogen  sulphide  precipitates  brown 
stannous  sulphide, 

SnCl2  +  H2S  <=±  SnS  +  2  HC1, 

which  after  standing  a  short  time  may  be  filtered  off. 

From  stannic  chloride,  or  metastannic  chloride,  hydrogen 

sulphide  forms  a  colloidal  precipitate  of  mixed  composition, 

partly  hydroxide  and  partly  sulphide: 

SnCl4  +  2  H2S  <=»  SnS2  +  4  HC1, 
SnCl4  +  3  H20  <=*  H2Sn03  +  4  HC1. 

This  precipitate  tends  to  assume  the  hydrosol  condition  in 
absence  of  coagulants;  but,  after  the  addition  of  ammonium 
nitrate  and  standing  to  settle,  the  precipitate  may  be  collected 
on  paper,  washed  with  an  acid  solution  of  ammonium  nitrate, 
and  dried.  The  precipitate  is  removed  from  the  paper  as  com- 
pletely as  possible  and  is  put  into  a  porcelain  crucible.  The 
paper  is  burned  in  a  platinum  spiral  and  the  ash  is  added  to  the 
precipitate  in  the  crucible.  Upon  ignition,  gently  at  first  to 
avoid  loss  of  stannic  sulphide  which  is  volatile  at  high  temper- 
ature, the  hydroxide  is  dehydrated  and  the  sulphide  is  converted 
to  oxide  containing  some  sulphate.  To  decompose  the  sulphate, 
the  crucible  is  cooled,  solid  ammonium  carbonate  (1  grm.)  is 
added,  the  cover  of  the  crucible  put  on,  and  the  ignition  re- 
peated. The  residue  is  weighed  as  stannic  oxide,  Sn02. 

The  Precipitation  of  Bismuth  Sulphide,  Bi2S3.  —  Bismuth 
sulphide,  precipitated  by  saturating  the  slightly  acid  solution 
with  hydrogen  sulphide,  may  be  collected  upon  asbestos  in  the 
perforated  crucible,  washed  with  water  charged  with  hydrogen 
sulphide,  freed  from  water  (which  repels  the  carbon  disulphide 
to  be  applied  later)  by  washing  with  alcohol,  extracted  with 
carbon  disulphide  to  remove  free  sulphur,  again  washed  with 
alcohol  and  with  ether,  dried  at  100  degrees,  and  weighed  as 
Bi2S3: 

2  Bi(N03)3  +  3  H2S  *±  Bi2S3  +  6  HN03. 

The  Precipitation  of  Lead  Sulphide,  PbS.  —  Lead  sulphide 
may  be  precipitated  completely  from  neutral  and  alkaline  solu- 


110  QUANTITATIVE  CHEMICAL  ANALYSIS 

tion  by  ammonium  sulphide  and  from  cold  and  slightly  acid 
solutions  by  hydrogen  sulphide: 

Pb(NO3)2  +  H2S  =  PbS  +  2  HN03. 

The  precipitate  is  prone,  however,  to  retain  some  of  the  acidic 
ion  of  the  soluble  salt  from  which  the  sulphide  is  formed,  and 
this  is  markedly  so  in  the  precipitation  from  solutions  con- 
taining chlorides.  The  process  is  generally  restricted  to  use  in 
separations,  other  means  for  the  determination  of  lead  being 
preferred. 

The  Precipitation  of  Cadmium  Sulphide,  CdS.  —  Cadmium 
sulphide  formed  by  precipitation,  like  lead  sulphide,  always  in- 
cludes products  of  the  partial  conversion  of  the  soluble  salt  to 
the  insoluble  sulphide,  such  as  Cd2Cl2S,  Cd2SO4S,  etc.,  and  is 
therefore  not  adapted  to  the  direct  determination  of  cadmium. 
However,  the  complex  precipitate  which  falls  when  hydrogen 
sulphide  is  passed  into  the  solution  of  the  cadmium  salt,  con- 
taining 2  to  7  per  cent  of  concentrated  sulphuric  acid  is  easily 
filterable,  readily  converted  to  cadmium  sulphate  by  evaporation 
with  sulphuric  acid,  and  may  be  brought  to  definite  form  for 
weighing  by  ignition  in  a  crucible  which  is  held  within  another 
crucible  so  that  the  walls  of  the  crucibles  are  1  cm.  apart. 
The  outer  crucible  may  be  heated  to  redness. 

The  Precipitation  of  Molybdenum  Sulphide,  MoS3.— 
Molybdenum  sulphide  may  be  precipitated  from  the  solution 
slightly  acid  with  sulphuric  acid  (or,  if  necessary,  with  hydro- 
chloric acid)  by  saturating  the  cold  solution  in  a  pressure  flask 
and  heating  the  closed  flask  on  the  water-bath  until  the  precipi- 
tate settles.  Or  the  ammoniacal  solution  of  the  molybdate 
may  be  charged  with  hydrogen  sulphide  until  it  is  bright  red 
and  then  acidified  to  bring  about  the  precipitation  of  the  sul- 
phide. The  precipitate  is  collected  on  asbestos  in  a  perforated 
crucible,  washed  with  acidulated  water  and  then  with  alcohol, 
and  dried  at  100  degrees. 

Sulphides  Precipitable  only  in  Alkaline  Solution 

Precipitation  of  Iron  Sulphide.  —  Iron  is  precipitated  as 
sulphide  when  it  is  desired  to  effect  a  separation  from  alumin- 
ium, chromium,  titanium,  and  uranium  which  are  not  precipi- 
table  as  sulphides  or  hydroxides  in  an  ammoniacal  tartrate 
solution.  The  acid  solution  containing  tartaric  acid  to  an 
amount  three  times  that  of  the  combined  oxides  is  saturated 


PROCEDURES  IN  GRAVIMETRIC  ANALYSIS  111 

with  hydrogen  sulphide,  and  then  made  just  alkaline  with 
ammonia  to  precipitate  ferrous  sulphide: 

FeS04  +  (NH4)2S  <=±  FeS  +  (NH4)2S04. 

After  settling,  the  precipitated  ferrous  sulphide  may  be  col- 
lected upon  a  filter,  washed  with  water  containing  a  little  am- 
monium sulphide,  dissolved  in  hydrochloric  acid,  and  oxidized 
with  nitric  acid  or  other  suitable  oxidizer.  From  the  solution 
of  the  oxidized  salt  ferric  hydroxide  may  be  precipitated  by 
ammonium  hydroxide. 

Precipitation  of  Manganous  Sulphide.  —  The  precipitation 
of  manganese  as  sulphide  is  a  useful  method  of  separating 
that  element  from  the  alkali-earth  elements,  magnesium,  cal- 
cium, strontium,  and  barium. 

If  the  amount  of  the  alkali-earth  elements  is  considerable 
the  precipitation  is  made  in  the  cold  solution.  To  the  slightly 
ammoniacal  solution,  contained  in  an  Erlenmeyer  flask  and 
charged  with  a  considerable  amount  of  ammonium  chloride 
or  nitrate,  a  slight  excess  of  freshly-made  ammonium  sulphide 
free  from  carbonates*  is  added;  and  the  flask,  nearly  filled 
with  previously  boiled  cold  water  and  stoppered,  is  set  aside 
for  at  least  twenty-four  hours;  the  clear  supernatant  liquid  is 
poured  off  upon  a  filter;  and  the  precipitate  is  washed  several 
times  by  decantation,  finally  transferred  to  the  same  filter,  and 
washed  with  water  containing  ammonium  nitrate  (5  per  cent) 
with  ammonium  sulphide  (1  per  cent). 

If  only  a  small  amount  of  the  alkali-earth  salts  is  present  the 
precipitation  may  be  made  by  adding  the  ammonium  sulphide 
to  the  boiling  solution  and  continuing  the  boiling  until  the  color 
of  the  precipitate  is  a  dirty  green.  After  settling  for  some 
minutes  the  precipitate  may  be  filtered  off  and  washed  with 
the  washing  mixture  described  above. 

The  precipitate  and  paper  are  dried  and  ignited  separately, 
the  paper  being  burned  in  a  spiral  of  platinum  wire.  The  ash 
and  the  precipitate  may  be  ignited  to  constant  weight  in  a 
crucible  well  enveloped  in  the  oxidizing  flame  of  a  powerful 

*  See  foot-note  on  p.  100. 


112  QUANTITATIVE  CHEMICAL  ANALYSIS 

burner,  and  weighed  as  manganous  manganite,  Mn304;  or  the 
residue  may  be  mixed  with  pure  sulphur,  heated  in  an  atmos- 
phere of  hydrogen,*  and  weighed  as  manganous  sulphide,  MnS. 
Precipitation  of  Nickel  Sulphide  and  Cobalt  Sulphide.  - 
Nickel  sulphide  and  cobalt  sulphide  may  be  precipitated  by 
ammonium  sulphide: 

NiCl2  +  (NH4)2S  <=»  NiS  +  2  NH4C1, 
CoCl2  +  (NH4)2S  <=»  CoS  +  2  NH4C1. 

In  this  process  (which  is  of  value  in  making  the  separation  of 
nickel  and  cobalt  together  from  alkali  and  alkali-earth  elements) 
the  neutral  solution  of  the  salts  (preferably  chlorides)  contain- 
ing ammonium  chloride  is  treated  in  a  flask  with  freshly  made 
ammonium  sulphide  free  from  carbonate,!  added  dropwise  until 
no  further  precipitation  takes  place,  and  the  flask  is  nearly 
filled  with  boiled  water.  After  standing  (12  hours),  the  pre- 
cipitate is  filtered  off  and  washed  with  a  solution  of  ammonium 
acetate  (5  per  cent)  containing  hydrogen  sulphide.  If  too 
much  ammonium  sulphide  is  used  in  the  precipitation  traces  of 
nickel  sulphide  will  dissolve  to  form  a  brown  solution  from 
which  the  nickel  sulphide  may  be  recovered  by  acidifying  with 
acetic  acid  and  heating.  In  any  event  the  filtrate  should  be 
acidified  with  acetic  acid,  heated  to  boiling,  and  treated  with 
hydrogen  sulphide  and  again  filtered  to  recover  traces  of  nickel 
sulphide  which  are  to  be  added  to  the  main  precipitate.  On 
ignition  the  precipitated  sulphides  of  nickel  and  cobalt  leave  a 
mixture  of  sulphides  and  oxide  which  may  be  dissolved  in 
aqua  regia,  converted  to  chlorides  by  repeated  evaporations 
with  hydrochloric  acid,  and  made  ready  for  the  separation  and 
determination  of  nickel  and  cobalt  by  appropriate  methods.  % 

*  See  p.  101.  t  See  foot-note  on  p.  100. 

J  The  metals  may  be  determined  together  electrolytically,  by  the  process 
described  on  p.  53  for  the  deposition  of  nickel.  For  the  separation  and  de- 
termination of  nickel  and  cobalt,  see  Tread  well-Hall :  Analytical  Chemistry, 
3rd.  ed.,  Vol.  II,  p.  161. 


CHAPTER   IV 
PROCEDURES  IN  VOLUMETRIC  ANALYSIS 

In  processes  of  Volumetric  Analysis  reactions  are  generally 
brought  about  by  adding  measured  volumes  of  reagent  solutions, 
made  to  contain  known  weights  of  reagents,  to  solutions  of  the  sub- 
stances to  be  determined;  and  the  amounts  of  the  substances  to  be 
determined  are  calculated  from  the  amounts  of  reagent  solutions 
employed.  The  process  of  measuring  out  the  reagent  solution 
is  known  as  titration.  According  to  the  kind  of  reaction  upon 
which  they  depend,  nearly  all  processes  of  volumetric  analysis 
may  be  classified  as  Neutralization  Processes,  Oxidation  Proc- 
esses, lodometric  Processes,  Precipitation  Processes,  Colorimetric 
Processes,  or  Gasometric  Processes. 

Normal  Solutions.  —  In  any  given  reaction  attention  is 
naturally  directed  to  the  function  of  particular  elements  or 
ions.  Thus,  in  reactions  involving  the  neutralization  of  acids, 
such  as  the  neutralization  of  hydrochloric  acid  by  an  alkali 
hydroxide, 

HC1  +  MOH  =  MCI  +  H20, 
or 

H  +  a  +  M  +  OH  =  M  +  Cl  +  H2O, 

attention  is  especially  directed  to  the  hydrogen  ion,  and  to 
the  hydroxyl-ion.  In  certain  reactions  of  oxidation,  such  as 
oxidations  by  potassium  permanganate  in  presence  of  sulphuric 
acid, 

2  KMn04  +  4  H2S04  =  2  KHS04  +  2  MnS04  =  3  H20  +  5  6, 

it  is  the  function  of  the  available  oxygen  which  is  prominent. 
In  some  processes  several  elements  or  ions  may  demand  con- 
sideration, as  in  the  iodometric  estimation  of  hydrochloric  acid 
by  the  determination  of  the  iodine  which  is  liberated  in  the 
action  of  that  acid  upon  a  mixture  of  potassium  iodide  and 

113 


114  QUANTITATIVE  CHEMICAL  ANALYSIS 

iodate,  the  free  iodine  being  determined  by  the  action  of  sodium 
arsenite  in  presence  of  an  alkali  acid  carbonate: 


2  12  +  As2O3  +  n(NaHCO3)  =  4  Nal  +  As205 

+  (n  -  4)NaHC03  +  2  H20  +  4  CO2. 

In  these  reactions  the  elements  or  ions  singled  out  for  special 
consideration  are  the  hydrogen  of  the  hydrochloric  acid,  the 
iodine  set  free,  and  the  oxygen  which  converts  the  arsenious 
oxide  (present  as  an  arsenite)  to  arsenic  oxide  (existing  as  an 
arsenate)  . 

In  order  to  arrange  the  concentrations  of  the  solutions  of 
hydrochloric  acid,  alkali  hydroxide,  potassium  permanganate, 
arsenic  trioxide  (in  the  arsenite),  and  iodine  (which  is  useful  in 
estimating  an  excess  of  the  arsenite  in  the  titration  process)  so 
that  they  may  correspond  to  one  another  volume  for  volume,  it 
is  necessary  to  take  into  consideration  both  the  gram-equivalent 
weights  of  these  reagents  and  the  reaction  valences  of  the 
elements  or  ions.  To  make  solutions  of  corresponding  con- 
centrations there  must  be  present,  in  equal  amounts  of  solution, 
amounts  of  the  reagents  proportional  to  the  quotients  of  their 
gram-equivalents  by  the  sum  of  the  active  valences  of  the 
specially  considered  element  or  ion.  For  example,  in  pre- 
paring the  reagents  mentioned  above  in  solutions  of  equal  con- 
centration in  respect  to  the  reactions  specified,  the  solutions 
are  so  made  that  a  given  volume  shall  contain  the  reagents  in 
amounts  proportionate  to  the  numerical  values  of  the  ex- 
pressions 

HC1     NaOH      Ba(OH)2     I      2  KMnQ4       KMnQ4    AsA 
T~      ~T~          ~^~     '   1>       ~10~  ~5~       ~^~ 

When  the  solutions  are  made  in  each  case  to  contain  in  1  liter 
a  weight  of  the  reagent  equal  to  the  quotient  obtained  by 
dividing  the  gram-equivalent  by  the  sum  of  the  utilized  va- 
lences of  the  specially  considered  element  or  ion  in  the  specified 
reaction  it  is  called  a  normal  solution. 

Standard  Solutions.  —  The  strengths  of  reagent-solutions 
are  conveniently  arranged  in  multiples  or  fractions  of  the 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS          115 

normal  strengths.  It  is  customary  to  make  use  of  normal, 
twice-normal,  half -normal,  fifth-normal,  tenth-normal,  etc., 
solutions,  according  to  circumstances,  indicating  these  solu- 
tions by  N,  2  N,  ^  or  N/2,  ^  or  N/5,  ^  or  N/10,  etc.;  and 

Z  o  L(J 

it  is  frequently  convenient  to  make  use  of  decimal  coefficients, 
as  in  1.8  N,  0.1  N,  0.52  N,  etc.,  it  being  much  easier  to  deter- 
mine the  exact  value  of  solutions  than  to  adjust  them  in  the 
ratio  of  even  numbers.  Solutions  of  determined  strengths  of 
whatever  value  are  called  standardized  solutions  or  standard 
solutions. 

Varying  Normality.  —  It  is  to  be  noted,  however,  that  a 
solution  which  is  normal,  or  which  bears  a  certain  relation  to 
normality,  for  one  sort  of  reaction  may  be  of  quite  different 
concentration  in  respect  to  a  reaction  which  proceeds  on  other 
lines.  For  example,  as  shown  above,  a  normal  solution  of 
potassium  permanganate  for  use  in  presence  of  sulphuric  acid 
contains  in  1  liter  the  number  of  grams  indicated  by  the  ex- 
pression -  — - — 4 .  When,  however,  oxidation  is  brought  about 
o 

by  the  permanganate  in  alkaline  solution,  as  in  Volhard's 
method  of  titrating  a  manganous  salt  in  presence  of  zinc  oxide, 

3  MnS04  +  2  KMnO4  +  7  ZnO  =  5  Zn02MnO  +  K2SO4+2  ZnS04, 

three  atoms  of  oxygen,  instead  of  the  five  atoms  which  enter 
into  the  reaction  of  oxidation  in  presence  of  acid,  are  avail- 
able from  the  2  KMn04  for  oxidizing  the  MnO  of  the  man- 
ganous salt  to  the  condition  of  MnO2.  For  reactions  of  this 
sort,  therefore,  the  normal  solution  of  permanganate  will  con- 
tain in  1  liter  a  weight  in  grams  of  permanganate  which  is 

.      2  KMnO4        KMnO4 
indicated  by  the  expression or 5 

O  o 

NEUTRALIZATION  PROCESSES 

(Acidimetry  and  Alkalimetry) 

The  Use  of  Indicators.  —  In  the  determination  of  an  acid 
by  means  of  a  neutralization  process,  a  standard  solution  of 


116  QUANTITATIVE  CHEMICAL  ANALYSIS 

base  is  required;  and  in  the  similar  determination  of  a  base  a 
standard  solution  of  acid  is  necessary.  In  such  determinations, 
of  either  sort,  the  condition  of  exact  neutralization,  the  end- 
point,  is  found  by  the  aid  of  the  indicator.  The  indicators  com- 
monly used  to  mark  the  end-point  in  processes  of  neutralization 
are  organic  substances,  the  colors  of  which  are  modified  by  the 
action  of  acids  and  bases,  and  of  the  numerous  indicators 
which  have  been  proposed,  methyl  orange  and  phenolphthalein 
are  the  most  widely  used  and  generally  available. 

Methyl  Orange  (Helianthin). —  Methyl  orange  gives  to  a 
solution  containing  free  alkali  a  yellow  color.  Upon  the  addi- 
tion of  an  excess  of  hydrochloric  acid  the  color  of  the  solution 
changes  to  red,  and  by  treatment  with  a  sufficient  amount  of 
alkali  hydroxide  the  yellow  color  is  restored. 

The  indicator  solution  is  prepared  by  treating  0.02  gr.ni.  of 
the  solid  reddish  violet  methyl  orange  (the  free  acid)  with  100 
cm.3  of  hot  water,  allowing  the  solution  to  cool,  and  filtering 
off  any  residue  (metasulphonic  acid);  or  0.022  grm.  of  the 
yellow  sodium  salt  may  be  treated  with  100  cm.3  of  hot  water 
containing  0.67  cm.3  of  N/10  hydrochloric  acid,  the  solution 
being  filtered  after  standing  some  time. 

The  color  of  alkaline  solutions  of  methyl  orange  is  due  to  the 
presence  of  yellow  di-methyl-amino-azo-benzene  sulphonate, 

I  MS03  -  </OH?)>  -  N  =  N  -  /QH^\  -  N(CH3)2, 

which  by  the  action  of  hydrochloric  acid  is  changed  to  di- 
methyl-amino-azo-benzene  sulphonic  acid,  also  yellow, 

II  HSO3- 

This  substance,  which  is  amphoteric,  possessing  the  acidic 
group  HS03  and  the  basic  group  N(CH3)2,  may,  according  to 
accepted  views,*  rearrange  to  the  isomeric  form  which  is  red, 

TT  

III  S03-<^C6H4  >  -  N  -  N  =  (^CeHO  =  N(CH3)2, 

*  Stieglitz:  Jour.  Am.  Chem.  Soc.,  26,  1112. 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS 


117 


and  this  with  hydrochloric  acid  forms  a  compound  (chloride 
and  sulphonic  acid)  also  red, 

H 

-N-N  = 


IV     HS03- 


C6H4  >  =  N(CH3)2C1. 


The  yellow  color  of  the  sulphonate  and  of  the  free  sulphonic 
acid  is  ascribed  to  the  presence  of  the  special  group  —  N  =  N  — 
known  as  the  azo-chromophor,  and  the  red  color  of  the  re- 
arranged compounds  is  attributed  to  the  presence  of  the  quinoid 
group  =  CeH4  =  which  is  known  as  the  quinoid-chromophpr. 

When  the  reddish  violet  scales  of  the  free  sulphonic  acid  are 
dissolved  in  the  least  possible  amount  of  water  the  solution 
takes  a  reddish  orange  color,  due  apparently  to  the  existence 
of  the  azo-form  (II),  and  the  quinoid  form  (III)  in  equilibrium: 


HS03- 


-  N(CHa)s 


=  N(CH3)2. 


On  dilution  with  water  the  solution  becomes  yellow,  the  azo 
form  predominating;  on  addition  of  a  trace  of  hydrochloric 
acid  the  solution  takes  the  red  color  due  to  the  quinoid  form 
and  its  salt  (IV).  Upon  further  dilution,  if  the  amount  of  acid 
has  not  been  excessive,  the  red  changes  again  to  yellow,  in  con- 
sequence of  the  hydrolytic  action  of  water  upon  the  quinoid 
salt  (IV)  and  the  reversion  of  the  quinoid  acid  (III)  thus  set 
free  to  the  azo  form  II. 

BEHAVIOR  OF  REPRESENTATIVE  ACIDS  AND  BASES 
TOWARD  METHYL  ORANGE 


Acids. 

Bases. 

Active. 

Inactive. 

Indeterminate. 

Active. 

Inactive. 

HC1 

HCN 

Organic 

NaOH 

C6H5NH2 

HNO3 

H2CO3 

acids  in 

KOH 

CH3  •  C6H4  •  NH2 

H2SO4 

H2S 

general. 

CaO2H2 

C6H5N 

H3P04* 

H3AsO3 

H3BO3 

Sr02H2 

etc. 

H2S03f 

BaO2H2 

MgO2H2 

NH4OH 

CH3NH2(H2O) 

C2H5NH2(H2O) 

*  Forming  H2MPO4. 


t  Forming  HMSO3. 


118  QUANTITATIVE  CHEMICAL  ANALYSIS 

Methyl  orange  is  a  sensitive  and  reliable  color  indicator  for 
a  few  of  the  stronger  acids  (showing  red);  and  it  is  a  good 
indicator  for  nearly  all  the  common  bases  (showing  yellow).  It 
may  therefore  be  used  to  indicate  the  mutual  neutralization  of 
such  acids  and  such  bases. 

Inasmuch,  however,  as  the  change  of  color  depends  not  only 
upon  the  relation  between  the  amounts  of  base  and  acid  present 
but  also  upon  the  concentrations  of  acid  and  indicator,  it  is 
essential  that  standardizations  and  titrations  should  be  made 
at  similar  concentrations  and  the  amount  of  the  indicator  used 
should  be  the  least  possible.  In  a  volume  of  100  cm.3  the 
change  in  color  is  distinct  when  titrations  are  made  with  fifth- 
normal  (or  stronger)  solutions,  while  with  tenth-normal  solu- 
tions the  change  at  the  point  of  neutrality  is  less  characteristic. 
A  few  drops  (five  or  six)  of  a  0.02  per  cent  solution  of  methyl 
orange  should  be  a  suitable  amount  of  the  indicator.  In 
titrating  an  acid  with  standard  alkali  the  end-point  of  the 
titration  should  be  determined  by  matching  the  color  of  the  solu- 
tion with  that  obtained  by  adding  a  similar  amount  of  the  in- 
dicator to  100  cm.3  of  water,  and  this  point  is  reached  when 
an  amount  of  alkali  exactly  equivalent  to  the  acid  has  been 
added;  but  in  the  reverse  process,  the  titration  of  an  alkali  by 
standard  acid,  a  slight  excess  of  acid  is  necessary  to  overcome 
the  hydrolytic  effect  of  the  water  upon  the  methyl-orange  salt. 
Theoretically,  the  results  should  be  more  accurate  when  acid 
is  titrated  by  alkali  than  when  alkali  is  titrated  by  acid.  Prac- 
tically, it  is  easier  for  most  observers  to  see  the  change  from 
yellow  to  red  rather  than  the  reverse  change  from  red  to 
yellow. 

Phenolphthalein.  —  The  indicator  solution  of  phenolphthalein 
is  prepared  by  dissolving  1  grm.  of  pure  phenolphthalein  in  100 
cm.3  of  86  per  cent  alcohol. 

Phenolphthalein  (I),  the  lactone  of  a  tribasic  acid,  is  color- 
less. From  it  may  be  prepared,  by  the  action  of  sodium 
hydroxide,  a  tribasic  salt  (IV)  which  is  colorless,  a  dibasic  salt 
(III)  which  is  red,  and  a  monobasic  salt  (II),  which  is  colorless. 
The  red  color  of  the  dibasic  salt  appears  to  be  associated  with 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          119 


chromophoric  group  =CeH4  =  .     The  relations  of  phenolphthal- 
ein  and  these  derivatives  are  shown  in  the  following  symbols.* 


>-O-H 


-0-H 


0  =  0- 


0 

Phenolphthalein 
(Colorless  lactone) 


=  C-0-Na 

Monobasic  salt 
(Colorless) 


CUD 


C-O-Na 

Dibasic  salt 
(Red) 


(IV) 


C-O-Na 

Tribasic  salt 
(Colorless) 


The  use  of  phenolphthalein  as  an  indicator  depends  upon  the 
formation  of  a  red  dibasic  salt  by  the  action  of  a  strong  base 
and  the  production  of  a  colorless  monobasic  salt  (or  the  lactone 

*  Kober  and  Marshall:  Jour.  Am.  Chem.  Soc.,  33,  59,  1780.    34,  1424. 


120  QUANTITATIVE  CHEMICAL  ANALYSIS 

of  the  free  acid)  by  the  action  of  an  acid.  It  is  to  be  noted, 
however,  that  the  red  color  of  the  dibasic  salt  which  is  produced 
when  an  alkali  hydroxide  is  added  to  a  solution  containing 
phenolphthalein  may  be  made  to  disappear  with  the  formation 
of  a  tribasic  salt  upon  excessive  addition  of  the  reagent  and  to 
return  when  the  dibasic  salt  is  reproduced  by  hydrolysis  on 
suitably  diluting  the  solution.  When  dissolved  in  water,  the 
colorless  monobasic  salt  which  is  formed  in  the  first  stage  of 
the  neutralization  by  the  base  is  also  slowly  acted  upon  by 
water  with  the  formation  of  colorless  phenolphthalein  and  the 
red  dibasic  salt  which  is  the  most  stable  of  these  salts  in  the 
water  solution.  It  is  obvious,  therefore,  that  standardizations 
and  titrations  should  be  executed  under  similar  conditions. 
Phenolphthalein  does  not  form  stable  salts  with  weak  bases  and 
is  not  available  in  the  titration  of  such  bases  nor  of  ammonia, 
but  is  a  good  indicator  for  strong  bases  and  for  many  adds, 
organic  and  inorganic.  It  is  even  sensitive  to  carbonic  acid, 
and  (at  0  degree  and  in  presence  of  sodium  chloride)  may  be 
used  to  indicate  the  point  at  which  neutral  sodium  carbonate 
is  converted  to  sodium  hydrogen  carbonate  (see  p.  127). 

The  Preparation  of  Standards 

Many  methods  have  been  proposed  for  the  standardization 
of  solutions  used  in  the  neutralization  processes.  For  most 
purposes  it  is  convenient  to  make  the  acid  the  primary  stand- 
ard and  to  fix  by  reference  to  it  the  standard  of  the  alkali. 
Hydrochloric  acid  in  solution  of  suitable  strength  makes  a  con- 
venient primary  standard  solution  for  use  with  either  methyl 
orange  or  phenolphthalein  as  the  indicator. 

The  Standard  Solution  of  Hydrochloric  Acid.  —  To  make  a 
solution  of  hydrochloric  acid  of  approximately  standard  strength 
it  is  only  necessary  to  dilute  the  concentrated  acid  suitably 
with  water.  If  pure  concentrated  hydrochloric  acid  is  diluted 
until  the  specific  gravity  of  the  solution  is  about  1.020  a 
nearly  normal  solution  is  obtained,  containing  approximately 
1  grm.  mol.  or  36.468  grm.  to  the  liter.  From  such  a  solution 
more  dilute  solutions  of  strength  suitable  for  use  in  titrimetric 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS 


121 


processes,  N/5  or  N/10,  may  be  made  by  properly  diluting 
measured  portions  with  water.  The  strength  of  such  solutions 
is  then  determined  exactly  by  precipitating  and  weighing  in 
the  manner  previously  described  for  the  gravimetric  determi- 
nation of  the  chloride  ion,*  the  silver  chloride  obtained  from 
an  exactly  measured  portion  of  the  solution. 

Experimental  Process:  Standardization  of  Hydrochloric  Acid.  —  From  an 
approximately  N/10  solution  of  hydrochloric  acid,  made  by  diluting 
10  cm.3  of  the  concentrated  acid  of  sp.  gr.  1.02  to  a  liter,  draw  from  a 
burette  a  100-cm.3  portion,  dilute  it  with  water  to  a  volume  of  300  cm.3, 
add  a  few  drops  of  nitric  acid  and  then  (gradually  and  with  constant 
stirring  to  coagulate  the  precipitate)  a  5  per  cent  solution  of  silver  nitrate, 
until  no  further  precipitation  takes  place.  Heat  the  liquid  to  boiling  and 
set  the  whole  aside  (in  the  dark)  to  cool.  When  completely  settled,  the 
precipitate  is  filtered  off  on  asbestos  in  a  perforated  crucible,  previously 
weighed,  washed  thoroughly  with  water  containing  a  few  drops  of  nitric  acid 
to  the  liter  and  then  with  a  little  pure  water.  Dry  first  at  100  degrees 
and  then  in  an  air-bath  heated  to  at  least  130  degrees  or  over  a  low  flame 
at  a  temperature  not  exceeding  the  fusing  point  of  silver  chloride.  From 
the  weight  of  silver  chloride  found  calculate  the  corresponding  weight  of 
hydrochloric  acid  and  the  weight  of  that  reagent  in  a  liter  of  the  solution 
analyzed.  State  the  results  in  the  following  form: 


Solution  of  HC1 
taken. 

AgCl  found. 

HC1 

Normality  of 
solution. 

corresponding  to 
AgCl  found. 

in  1  liter  of 
solution. 

Cm'. 

Grm. 

Grm. 

Grm. 

The  Standard  Solution  of  Sodium  Hydroxide.  —  A.  It  is 

not  necessary  that  sodium  hydroxide  should  be  entirely  free 
from  carbonate  if  the  solution  is  to  be  used  in  the  neutralization 
of  hydrochloric  acid  (or  other  strong  acid)  with  methyl  orange 
as  the  indicator.  For  this  purpose  the  solution  may  be  prepared 
by  washing  the  surface  of  a  suitable  amount  of  sodium  hydrox- 
ide, sticks  or  lumps,  with  water  (to  remove  the  outer  coating 

*  See  p.  67. 


122  QUANTITATIVE  CHEMICAL  ANALYSIS 

of  carbonate)  and  dissolving  the  washed  hydroxide  in  water. 
To  make  a  nearly  normal  solution,  about  1  grm.  mol.  of  sodium 
hydroxide  (40  grm.)  is  dissolved  in  a  liter  of  water,  and  more 
dilute  solutions,  N/5  or  N/10,  may  be  made  by  diluting  with 
water.  The  necessary  degree  of  dilution  may  be  fixed  ap- 
proximately by  titrating  a  small  portion,  roughly  measured, 
against  the  standardized  hydrochloric  acid,  with  methyl  orange 
as  the  indicator,  and  the  exact  standard  of  the  diluted  solution 
may  then  be  fixed  by  a  similar  titration  carefully  executed. 

Experimental  Process:  Standardization  of  NaOH.  —  Weigh  out  ap- 
proximately 10  grm.  of  sodium  hydroxide  in  sticks,  wash  away  the  super- 
ficial coating  of  carbonate,  and  dissolve  the  rest  in  200  cm.3  of  water. 

A.  Measure  from  a  burette  4  cm.3  of  the  strong  solution  just  made, 
add  a  few  drops  of  the  methyl  orange  indicator,  and  run  in  N/10  hydro- 
chloric acid  to  the  point  of  neutralization.     From  the  result  calculate 
how  much  of  the  strong  solution  of  sodium  hydroxide  must  be  taken  to 
make  an  N/10  solution  (40.01  grm.  to  the  liter)  and  make  up  a  suitable 
amount  of  an  approximately  N/10  solution. 

B.  Draw  from  a  burette  40  cm.3  of  the  approximately  N/10  solution 
of  sodium  hydroxide,  carefully  mixed  and  at  the  laboratory  temperature, 
add  a  few  drops  of  the  methyl  orange  indicator  and  titrate  to  the  end- 
point  with  N/10  hydrochloric  acid.     Calculate  the  exact  value  of  the 
sodium  hydroxide  solution  and  label  it  in  terms  of  "value  per  cm.3"  and 
in  terms  of  normality. 

B.  For  titration  against  acids  in  presence  of  phenolphthalein 
used  as  the  indicator,  the  standard  solution  of  sodium  hydroxide 
must  be  prepared  free  from  carbonate,  since  the  carbon  dioxide 
set  free  in  the  process  of  neutralization  is  capable  of  converting 
the  red  dibasic  sodium  salt  into  the  colorless  monobasic  sodium 
salt,  thus  vitiating  the  apparent  balance  between  the  alkali 
hydroxide  and  the  standard  acid. 

For  use  with  phenolphthalein  the  N/10  solution  is  most 
convenient  and  accurate.  Such  a  solution  of  sodium  hydroxide, 
free  from  carbonate,  is  best  prepared  by  the  action  of  water 
vapor  upon  clean  metallic  sodium  in  absence  of  carbon  dioxide. 

A  device  for  this  purpose  (Kuster)  is  shown  in  Fig.  23.  A 
tubulated  bell-jar  sealed  at  the  bottom  by  water  in  a  flat- 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS 


123 


bottomed  dish  and  standing  upon  a  glass  triangle,  to  give  a 
vent  through  the  water  to  the  hydrogen  evolved  in  the  reaction, 
makes  a  suitable  isolating  chamber.  A  cone  of  nickel  gauze,, 
suspended  over  the  nickel  vessel  which  is  to  receive  the  sodium 
hydroxide  as  it  is  formed,  serves  to  contain  the  sodium.  In 
using  this  apparatus,  a  suitable  amount  of  sodium,  freshly  cut 
to  give  a  clean  surface,  is  put  in  the  gauze  cone  and  the 
funnel,  the  dish  containing  water,  and  the  vessel  to  receive  the 
sodium  hydroxide  are  arranged  as  shown.  As  the  sodium  is 


FIG.  23. 


FIG.  24. 


attacked  by  the  water  vapor  sodium  hydroxide  is  formed  and, 
dissolving  in  the  vapor,  drops  into  the  vessel  below.  The 
operation  may  be  hastened  by  gently  heating  the  water  in  the 
dish  upon  a  steam  radiator,  after  the  air  has  been  replaced  by 
the  hydrogen.  The  solution  of  sodium  hydroxide  in  the  con- 
taining vessel  is  suitably  diluted  with  distilled  water,  free  from 
carbon  dioxide,  so  that  a  liter  of  the  solution  shall  contain  the 
hydroxide  produced  by  2.3  grm.  of  sodium,  and  is  transferred 
as  quickly  as  possible  to  a  bottle  provided,  as  shown  in  Fig.  24, 


124 


QUANTITATIVE  CHEMICAL  ANALYSIS 


with  a  protecting  tube  filled  with  soda-lime,  a  side-neck  burette 
fitted  with  a  pinch-cock  or  the  Bunsen  bead-tap,*  and  a  siphon- 
tube  for  filling  the  burette.  The  exact  standard  of  the  solution 
is  fixed  by  titration  with  the  standard  acid.f 

The  Standard  Solution  of  Barium  Hydroxide.  —  For  many 
purposes,  the  more  easily  prepared  solution  of  barium  hydroxide 
may  be  substituted  for  the  solution  of  sodium  hydroxide  free 
from  carbonate.  Such  a  solution  may  be  easily  made  by  dis- 
solving crystallized  barium  hydroxide  in  a  suitable  amount  of 
water,  diluting  to  the  volume  desired,  allowing  the  solution  to 
stand  until  all  barium  carbonate  has  settled  out,  and  siphoning 
into  a  bottle  filled  with  air  free  from  carbon  dioxide,  protected 
by  a  soda-lime  tube,  and  provided  with  burette  and  siphon,  as 
shown  in  Fig.  24.  The  exact  standard  of  the 
solution  may  be  determined  by  titration*  with 
standard  hydrochloric  acid. 

Experimental  Process:  Standardization  of  Barium 
Hydroxide,  with  Phenolphthalein.  —  Weigh  out  about 
20  grm.  of  crystallized  barium  hydroxide  and  dissolve 
it  in  a  suitable  amount  of  water  contained  in  a  closed 
flask.  Let  the  solution  stand  for  a  day  or  two,  until 
the  precipitated  barium  carbonate  has  settled  com- 
pletely. Fill  the  bottle  to  be  used  as  a  container 
with  water  and  displace  the  water  by  air  freed  from 
carbon  dioxide.  Draw  the  solution  of  barium  hy- 
droxide through  a  siphon  filter  (made  from  a  perforated 
crucible  with  the  asbestos  filter,  stopper,  and  tube,  as  shown  in  Fig.  25) 
into  the  container,  adjust  the  protecting  tube,  siphon,  and  burette,  as 
shown  in  Fig.  24.  To  fix  the  standard,  a  50-cm.3  portion  of  N/10 
hydrochloric  acid  is  placed  in  a  flask,  a  drop  of  phenolphthalein  is 
added,  and  barium  hydroxide  is  run  in  from  the  burette  until  the  red  color 
of  the  dibasic  salt  appears.  Calculate  the  strength  of  the  solution  in  terms 
of  standard  hydrochloric  acid  and  state  the  results  in  the  following 
form: 

*  To  draw  liquid  from  the  burette  the  tube  is  held  between  the  thumb 
and  forefinger  over  the  bead  and  by  gentle  pressure  the  tube  is  deformed 
so  that  the  liquid  may  flow. 

t  See  Experimental  Process,  p.  122. 


FIG.  25. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS 


125 


HC1  solution 
taken. 

Ba(OH)2 
solution  used. 

Ba(OH)2  solution 
corresponding  to 
1  cm.3  of  HC1 
solution. 

Ba(OH)2  in  1  cm.3 
of  solution. 

Normality  of 
Ba(OH)2 
solution. 

Cm." 

Cm.3 

Cm.3 

Grm. 

Experimental  Process:  Determination  of  Stronger  Acids  in  Dilute  Solu- 
tion. —  Measure  out  roughly  25  cm.3  of  N/10  sodium  hydroxide  or  barium 
hydroxide,  introduce  methyl  orange,  and,  from  a  pipette  graduated  to 
tenths  of  a  cubic  centimeter,  add  portions  of  the  dilute  hydrochloric  acid, 
nitric  acid,  or  sulphuric  acid  of  unknown  strength  (e.g.,  the  desk  supply) 
until  the  end-point  is  reached.  Then  from  the  results  of  this  rough  com- 
parison calculate  the  volume  of  the  acid  necessary  to  make  100  cm.3  of  acid 
of  approximately  N/10  concentration.  Measure  this  calculated  volume 
of  the  original  acid  into  a  counterpoised  100-cm.3  graduated  flask,  weigh 
carefully,  and  dilute  to  the  mark. 

A.  Measure  from  a  burette  40  cm.3  of  the  prepared  solution.  Add 
methyl  orange  and  titrate  to  the  end-point  with  standard  sodium  hydrox- 
ide, approximately  N/10.  From  the  amount  used  of  the  sodium  hydroxide 
solution  calculate  the  percentage  strength  (by  weight)  of  the  original  acid. 
Tabulate  the  results  in  the  form  given  below. 


Acid  used. 

Weight  of 
original 
acid  in  100 
cm.3  of  the 
prepared 
solution. 

Volume  of 
prepared 
solution 
taken. 

Standard 
solution 
of  alkali 
used. 

Weight  of 
NaOH  or 
Ba(OH)2 
used. 

Weight  of  acid 
equivalent  to 
NaOH  or 
Ba(OH)2used. 

Concen- 
tration of 
original 
acid  by 
weight. 

Grm. 

Cm.3 

Cm.3 

•      Grm. 

Grm. 

Per  cent. 

Per  cent. 

HC1 

HNO3 

H2SO4 

HC1 

HNO3 

H2SO4 

126  QUANTITATIVE  CHEMICAL  ANALYSIS 

B.  Measure  from  a  burette  40  cm.3  of  the  prepared  solution  made  to 
contain  a  known  weight  of  the  original  acid,  add  phenolphthalein,  and 
titrate  to  the  end-point  with  standard  barium  hydroxide,  approximately 
N/10.  From  the  amount  used  of  the  standard  barium  hydroxide  solution 
calculate  the  percentage  strength  (by  weight)  of  the  original  acid. 

Experimental  Process:  Determination  of  Sodium  Hydroxide  and  So- 
dium Carbonate  in  Crude  Caustic  Soda.  — Weigh  out  exactly  about  5  grm. 
of  crude  caustic  soda,  dissolve  it  in  water,  and  make  up  the  solution  to  1 
liter  at  the  atmospheric  temperature. 

Process  I.  —  A.  Of  this  solution,  draw  50  cm.3  from  a  burette,  add  five 
or  six  drops  of  methyl  orange,  and  add  nearly  N/10  hydrochloric  acid  to 
apparent  neutralization.  Reactions  take  place  according  to  the  equations 

NaOH  +  HC1  =     NaCl  +  H2O, 
Na2CO3  +  2  HC1  =  2  NaCl  +  H2O  +  CO2. 

With  standard  solutions  stronger  than  fifth-normal  the  dissolved  carbon 
dioxide  does  not  affect  the  end-point  with  methyl  orange;  but  with  tenth- 
'  normal  solutions  the  end-point  is  somewhat  hastened  by  the  carbon  dioxide. 
To  obviate  this  error,  it  is  only  necessary  to  expel  the  carbon  dioxide  by 
boiling,  to  cool  the  solution  (now  yellow)  and  to  add  the  slight  amount  of 
standard  hydrochloric  acid  necessary  to  produce  the  end-color.  The 
amount  of  hydrochloric  acid  used  is  the  measure  of  the  sodium  present  as 
hydroxide  and  as  carbonate. 

B.  Draw  50  cm.3  of  the  solution  from  the  burette,  add  more  than 
enough  barium  chloride  in  solution  to  affect  the  transformations  shown  in 
the  following  equations: 

NasCOs  +  BaCl2  =  2  NaCl  +  BaCO3  (insoluble), 
2  NaOH  +  BaCl2  *±  2  NaCl  +  Ba(OH)2  (soluble). 

Add  phenolphthalein  to  the  mixture  and  titrate  by  standard  hydrochloric 
acid  very  slowly  and  with  constant  stirring;  for,  unless  proper  precaution 
is  taken  in  respect  to  the  gradual  addition  of  the  hydrochloric  acid  some  of 
the  insoluble  barium  carbonate  may  be  attacked  before  the  indicator 
shows  the  end-point.  If  the  titration  is  properly  conducted,  the  hydro- 
chloric acid  should  be  used  up  in  the  reactions 

NaOH  +  HC1  =  NaCl  +  H2O, 
Ba(OH)2  +  2  HC1  =  BaCl2  +  2  H2O. 

The  amount  of  acid  used  in  B  measures  the  amount  of  sodium  hydroxide 
originally  present.  The  amount  of  acid  used  in  A  measures  the  amount  of 
sodium  hydroxide  and  sodium  carbonate  present.  The  difference  between 
the  amounts  of  acid  used  in  A  and  B  measures  the  sodium  carbonate 
originally  present. 

State  the  results  in  the  following  form: 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          127 


Solution 
NaOH+ 

Na2C03 
taken. 

Standard 
HC1  used 
for  com- 
plete neu- 
tralization. 
(A) 

Standard 
HC1  used 
in  neutral- 
izing the 
hydroxide. 
(B) 

Standard 
HC1  used 
in  neutral- 
izing car- 
bonate. 
(A-B) 

NaOH  in 
50  cm.* 
of  solu- 
tion. 

Na2C03 
in  50  cm.3 
of  solu- 
tion. 

In  crude  soda. 

NaOH 

Na2CO3 
Per  cent. 

Cm.3 

Cm.3 

Cm.3 

Cm.3 

Grm. 

Grm. 

Per  cent. 

Process  II.  —  A.   Draw  from  a  burette  50  cm.3  of  the  solution.     Add 
5  grm.  of  sodium  chloride  and  chill  the  solution  to  0  degree.     Add  phenol- 
phthalein  and  titrate  with  N/10  hydrochloric  acid  to  colorlessness: 
nNaOH  +  Na2CO3  +  (n  +  1)HC1  =  (n  +  l)NaCl  +  NaHCO3  +  nH20. 

B.  To  the  solution  just  titrated  add  methyl  orange  and  titrate  with  N/10 
hydrochloric  acid  to  the  end-color  (brownish  red) : 

NaHCO3  +  HC1  =  NaCl  +  H2O  +  CO2. 

Twice  the  amount  of  hydrochloric  acid  used  in  the  methyl  orange 
titration  (B)  will  measure  the  sodium  carbonate;  and  the  difference  be- 
tween the  amount  of  acid  used  in  the  phenolphthalein  titration  (A)  and 
that  used  in  the  methyl  orange  titration  (B)  will  measure  the  sodium 
hydroxide.  State  the  results  in  the  following  form: 


Solution  of 
nNaOH 
+Na2CO3 
taken. 

Standard  HC1 
used  with 
phenolphthal- 
ein. 

(A) 

Standard 
HC1  used 
with 
methyl 
orange. 
(*) 

Na2C03 
indicated. 

NaOH 
indicated. 

In  crude  soda. 

Na2C03 

NaOH 

Cm.3 

Cm.3 

Cm.3 

Grm. 

Grm. 

Per  cent. 

Per  cent. 

The  Determination  of  Boric  Acid.  —  Many  of  the  ordinary 
indicators,  even  methyl  orange,  are  more  or  less  sensitive  to 
the  action  of  boric  acid  in  fairly  concentrated  solutions  and  fail, 
therefore,  to  indicate  with  accuracy  the  exact  point  at  which  a 
stronger  acid  in  presence  of  free  boric  acid  may  be  removed  by 
neutralization  with  a  standard  alkali.  The  mixture  of  an 
iodide  with  an  iodate,  sensitive  to  the  action  of  stronger  acids, 


128  QUANTITATIVE  CHEMICAL  ANALYSIS 

is  not  acted  upon  by  boric  acid  when  present  to  the  extent  of 
the  equivalent  of  0.1  grm.  of  boric  anhydride  in  25  cm.3  of 
solution.  When  a  solution  containing  boric  acid,  of  suitable 
concentration,  and  hydrochloric  acid  is  treated  with  an  iodide- 
iodate  mixture,  and  the  liberated  iodine  taken  up  by  sodium 
thiosulphate,  the  solution  will  contain  only  neutral  chloride, 
neutral  iodide,  tetrathionate,  possibly  a  slight  excess  of  thio- 
sulphate, and  free  boric  acid.  Carbonic  acid,  which  if  present 
would  interfere  with  the  subsequent  use  of  phenolphthalein  as 
an  indicator,  may  be  removed  as  barium  carbonate  by  the 
addition  of  barium  chloride  in  presence  of  the  iodide-iodate 
mixture: 

5  KI  +  KI03  +  4  HC1  +  BaCl2  +  C02  +  H303B 

=  BaC03  +  6  KC1  +  H303B  +  2  H20  +  3  I2. 

When  boric  acid  and  mannite  (or  glycerine  in  suitably  Targe 
proportion)  are  mixed  in  solution,  a  peculiar  compound  of 
strongly  acid  properties  is  the  result.  This  compound  de- 
composes carbonates,  its  acid  taste  (comparable  to  that  of 
citric  acid)  is  much  stronger  than  that  of  boric  acid  alone,  and 
the  solution  shows  greater  electrical  conductivity  and  a  lower 
freezing  point  than  a  similar  molecular  solution  of  either  sub- 
stance alone.  Other  polyatomic  alcohols  (but  all  to  a  less  de- 
gree than  mannite)  and  some  organic  acids  show  this  peculiar 
property  of  combining  chemically  with  boric  acid  to  increase  its 
acid  qualities.  If,  therefore,  mannite  is  added  to  the  solution 
of  boric  acid  after  treatment  with  the  iodide-iodate  mixture 
and  thiosulphate,  and  with  barium  chloride,  the  boromannite 
compound  of  strongly  acid  properties  is  formed,  and  this  may 
be  titrated  with  an  alkali  hydroxide  in  presence  of  phenol- 
phthalein. 

In  the  practical  procedure,  the  solution  in  which  boric  acid 
is  present  to  an  amount  not  exceeding  0.1  grm.  in  25  cm.3  is 
made  slightly  acid  to  litmus  by  hydrochloric  acid  and  is  treated 
with  5  cm.3  of  a  10  per  cent  solution  of  barium  chloride.  To 
the  iodide-iodate  mixture  just  bleached  (as  is  generally  neces- 
sary) by  a  dilute  solution  of  sodium  thiosulphate  a  single  drop 
of  the  solution  to  be  analyzed  is  added.  If  the  iodine  color 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          129 

develops,  the  solution  is  known  to  be  acidic  with  hydrochloric 
acid  and  all  the  boric  acid  is  in  free  condition.  If  no  color  de- 
velops the  borate  solution  must  be  carefully  acidified  with 
hydrochloric  acid  until  a  drop  of  the  solution  will  liberate 
iodine  from  the  iodide-iodate  mixture.  Then,  the  iodide-iodate 
mixture,  containing  about  2  grm.  of  potassium  iodide  with 
0.5  grm.  of  potassium  iodate,  is  added  to  the  solution  of  the 
boric  acid,  and  the  liberated  iodine  is  bleached  with  sodium 
thiosulphate.  A  few  drops  of  the  phenolphthalein  indicator 
are  now  added,  and  N/5  sodium  hydroxide  is  run  in  until  a 
strong  red  coloration  is  produced.  The  addition  of  a  pinch  of 
mannite,  1  grm.  or  2  grm.,  has  the  effect  of  bleaching  the  color 
and  the  alkali  solution  is  run  in  again  until  a  faint  color,  per- 
manent on  the  addition  of  more  mannite,  is  obtained. 

Neither  silica  nor  fluorides  will  interfere  with  the  correct- 
ness of  this  process.  The  best  results  are  obtained  in  cold 
solution  of  final  volumes  not  greater  than  50  cm.3.  * 

Other  Standards  and  Indicators.  —  Other  acids  —  such  as 
nitric  acid,  sulphuric  acid,  and  oxalic  acid  —  are  sometimes 
used  with  advantage  as  standard  solutions  in  neutralization 
processes,  and  these  solutions  are  best  standardized  by  titra- 
tion  against  a  standard  alkali  solution,  itself  standardized 
against  standard  hydrochloric  acid  or  against  weighed  portions 
of  pure  succinic  acid,  H^C^^,  dried  in  air  or  over  sulphuric 
acid,  with  phenolphthalein  as  the  indicator. 

In  certain  cases,  standard  solutions  of  potassium  hydroxide 
or  of  ammonia  may  advantageously  replace  the  more  usual 
standard  solutions  of  sodium  hydroxide  or  barium  hydroxide. 

Other  indicators  are  sometimes  substituted  for  methyl  orange 
and  phenolphthalein. 

OXIDATION   PROCESSES 

Oxidation  processes,  or  processes  of  oxidimetry,  present  three 
distinct  phases  of  procedure: 

(/)  The  substance  to  be  determined  may  be  oxidized  definitely 
by  direct  titration  with  a  standard  oxidizer.  (II)  The  substance 
may  be  first  reduced  to  a  definite  condition  and  then  reoxidized 
*  Jones:  Methods  in  Analysis,  Gooch,  p.  206. 


130  QUANTITATIVE  CHEMICAL  ANALYSIS 

definitely  by  the  standard  oxidizer.  (Ill)  The  substance  may  be 
reduced  to  an  inert  form  by  a  definite  amount  of  a  reagent  taken  in 
an  excess  which  is  then  titrated  with  the  standard  oxidizer  to  find 
(by  difference)  a  measure  of  the  substance  reduced. 

Use  of  Potassium  Permanganate.  —  The  most  generally 
applicable  reagent  containing  oxygen  directly  available  for 
oxidizing  purposes  in  volumetric  analysis  is  potassium  per- 
manganate in  solutions  of  suitable  concentrations  —  commonly 
tenth-normal.  Solutions  made  by  dissolving  the  crystallized 
potassium  permanganate  in  water,  especially  if  hot  water  is 
used  to  hasten  the  process  of  going  into  solution,  invariably 
contain  in  suspension  manganese  dioxide  or  an  insoluble  potas- 
sium manganite.  If  the  higher  manganese  oxide,  or  the  man- 
ganite,  is  removed  by  filtering  the  solution  through  asbestos  * 
(paper  will  produce  further  reduction  of  the  permanganate) 
such  solutions  may  remain  essentially  unchanged  for  months, 
provided  they  are  kept  free  from  dust  and  are  not  exposed  to 
the  action  of  direct  sunlight.  The  standard  of  such  solutions 
should,  nevertheless,  be  redetermined  from  time  to  time.  In 
presence  of  a  suitable  acid,  such  as  sulphuric  acid,  potassium 
permanganate  acts  as  an  oxidizer  according  to  the  equation 

2  KMnO4  -h  4  H2SO4  =  2  KHSO4  +  2  MnSO4  +  3  H2O  +  5  6, 

and,  in  a  solution  devoid  of  free  acid,  in  the  sense  of  the  ex- 
pression 

2  KMn04  +  H20  =  2  MnO2  +  2  KOH  +  3  6. 

For  most  purposes  of  volumetric  analysis  potassium  perman- 
ganate is  made  to  react  in  presence  of  acid  and  a  solution  of 
tenth-normal  strength  is  generally  suitable.  Two  gram-mole- 
cules of  potassium  permanganate  liberate  5  gram-atoms  of 
oxygen,  each  atom  having  a  valence  of  two,  in  reactions  which 
take  place  in  presence  of  acid.  A  normal  solution  will  there- 

*  A  simple  form  of  asbestos  filter  may  be  made  by  plugging  a  funnel 
with  glass  wool  and  covering  the  latter  with  an  emulsion  of  asbestos  such 
as  is  used  in  the  perforated  crucible. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          131 

fore  contain  in  1  liter  that  weight  of  the  permanganate  which 
equals  two  gram-molecules  divided  by  ten: 

2.KMnO4      KMnO4      158.03 


10 


=  31.61  grm. 


A  solution  of  N/10  concentration  for  use  in  acid  solution  is 
therefore  made  to  contain  3.161  grm.  in  the  liter,  or  ^V  gram- 
molecule  of  potassium  permanganate. 

Potassium  permanganate  is  exceedingly  sensitive  to  the 
action  of  the  organic  matter  almost  always  present  in  water 
even  when  it  has  been  distilled,  and  the  oxide  of  manganese 
produced  in  such  action  behaves  as  a  catalyzer  in  promoting 
further  decomposition  of  the  permanganate.  In  making  an  ap- 
proximately N/10  solution  it  is  customary  to  weigh  out  about 
3.16  grm.  of  potassium  permanganate,  dissolve  it  in  1  liter  of 
hot  distilled  water,  filter  on  an  asbestos  filter,  set  aside  for 
several  days  and  then  filter  again. 

Standardization  of  Potassium  Permanganate.  —  Many 
modes  of  standardizing  the  nearly  N/10  solution  of  potassium 
permanganate  have  been  proposed  and  applied,  but  the  most 
convenient  method  is  based  upon  the  use  of  sodium  oxalate 
which  may  now  be  obtained  of  certified  purity.*  The  principle 
upon  which  this  process  of  standardization  depends  is  the 
oxidation  of  the  oxalic  acid  liberated  by  the  action  of  sulphuric 
acid  upon  the  sodium  oxalate. 

2  KMn04  +  5  Na2C204  +  8  HaSO* 

=  K2SO4  +  2  MnSO4  +  5  Na2SO4  +  5  H2C2O4  +  3  H2O  +  5  O 
=  K2SO4  +  2  MnS04  +  5  Na2SO4  +  10  C02  +  8  H20. 

A  weighed  amount  of  sodium  oxalate  is  dissolved  in  water 
heated  to  80  or  90  degrees,  sulphuric  acid  is  added,  and  the  hot 
solution  is  titrated  with  the  nearly  N/10  permanganate  solution. 
At  first  the  action  is  slow  and  the  permanganate,  contained  in 
a  glass-stoppered  burette,  should  be  added  so  gradually  that  the 
color  imparted  by  each  drop  shall  disappear  before  the  addition 
of  the  next.  After  a  small  amount  of  manganous  sulphate  has 

*  Purchaseable  from  the  Bureau  of  Standards,  Washington,  D.  C. 


132 


QUANTITATIVE  CHEMICAL  ANALYSIS 


been  produced,  the  speed  of  the  reaction  is  very  much  increased 
(by  the  catalytic  action  of  this  substance)  and  the  permanganate 
may  then  be  run  in  with  rapidity.  From  the  weight  of  sodium 
oxalate  taken  and  the  amount  of  the  permanganate  solution 
used  the  "  oxygen  equivalent  per  cm.3"  of  the  latter  may  be 
calculated. 

Experimental  Process:  Preparation  and  Standardization  of  N/10  Per- 
manganate. —  A.  Make  a  solution  of  potassium  permanganate  (of  ap- 
proximately 1/50  gram-molecular  strength  in  1  liter;  N/10  for  use  in 
presence  of  acid)  by  dissolving  3.16  grm.  of  potassium  permanganate  in 
hot  water,  filtering  on  an  asbestos  filter,  setting  aside  for  some  days,  re- 
filtering  into  a  liter  flask,  and  making  up  to  the  1  liter  mark. 

B.  Dissolve  about  0.25  grm.  of  sodium  oxalate  in  250  cm.3  of  water 
heated  to  80  or  90  degrees  and  contained  in  a  400-cm.3  beaker.  Add  10 
cm.3  of  [1  :  1]  sulphuric  acid.  Titrate  immediately  with  the  nearly  N/10 
solution  of  potassium  permanganate,  stirring  continually  and  taking  care 
not  to  run  in  the  permanganate  at  a  rate  greater  than  10  cm.3  to  15  cm.3 
per  minute  and  to  add  the  last  1  cm.3  drop  wise,  allowing  time  for  full 
decolorization  between  successive  drops  until  the  end-color  (which  should 
match  that  given  by  a  single  drop  of  the  permanganate  solution  to  300 
cm.3  of  water  containing  10  cm.3  of  (1  :  1)  sulphuric  acid)  remains.  The 
reaction  proceeds  according  to  the  equation 

4  +  H2S04  +  O  =  N*SO4  +  2  CO2  +  H2O. 


From  the  weight  (w)  of  sodium  oxalate  taken  and  the  number  of  cubic 
centimeters  (n)  of  the  solution  of  potassium  permanganate  employed  in 
oxidizing  the  oxalic  acid  produced  to  carbon  dioxide  and  water,  calculate 
weight  of  oxygen  used  and  the  "  oxygen  value"  per  cubic  centimeter  of 
the  permanganate. 

Na2C2O4  :  O  =  134  :  16  =  w  :  x  =  oxygen  used. 

-  =  "oxygen  value  per  cm.3.  " 
State  the  results  in  the  following  form: 


Na2C2O4  taken. 

Permanganate 
solution  used. 

Oxygen  employed  in 
oxidation. 

"  Oxygen  value  per 
cm.3  "  of  permanganate. 

Grm. 

Cm." 

Grm. 

Grm. 

PROCEDURES  IN   VOLUMETRIC  ANALYSIS          133 

Other  Modes  of  Standardizing  Potassium  Permanganate. — 

Solutions  of  potassium  permanganate  may  sometimes  be  ex- 
actly standardized  to  advantage  by  titration  against  a  ferrous 
salt  made  from  a  weighed  amount  of  electrolytic  iron,*  or 
against  a  weighed  amount  of  arsenic  trioxide,f  or  by  absorbing 
and  weighing  with  molecular  silver  the  iodine  set  free  by  the 
permanganate  from  potassium  iodide  in  an  acidified  solution.}: 

Direct  Oxidations  by  the  Permanganate  Solution 

Many  useful  analytical  processes  depend  upon  the  use  of  the 
standard  solution  of  potassium  permanganate  for  the  direct  oxi- 
dation of  the  substance  to  be  determined:  (7)  in  presence  of  free 
acid;  (II)  in  absence  of  free  acid.  Examples  of  such  processes 
are  found  in  (I)  the  determination  of  oxalates,  iron  in  ferrous 
salts,  and  hydrogen  dioxide  by  titration  in  presence  of  dilute 
acid;  and  of  nitrous  acid  in  concentrated  sulphuric  acid;  and 
(II)  the  determination  of  manganese  and  formates  by  titration 
in  absence  of  free  acid. 

Determination  of  Oxalates.  —  By  the  reversal  of  the 
process  by  which  potassium  permanganate  is  standardized  by 
means  of  sodium  oxalate  various  oxalates  may  be  determined. 
Moreover,  many  elements  capable  of  yielding  by  precipitation 
oxalates  of  definite  constitution  may  be  thus  estimated,  either 
by  direct  titration  with  standard  permanganate  or,  indirectly, 
by  titration  of  the  excess  remaining  from  a  definite  amount  of 
oxalic  acid  or  a  soluble  oxalate  used  to  precipitate  them.  Ele- 
ments which  may  be  estimated  in  this  manner  are  calcium, 
strontium,  barium,  copper,  lead,  zinc,  mercury  (in  the  mercurous 
form),  cerium,  and  lanthanum. § 

Experimental  Process:  Determination  of  Copper.  —  Weigh  out  carefully 
about  0.5  grm.  of  pure,  crystallized  copper  sulphate,  CuSO4  •  5  H2O,  and, 
dissolve  it  in  50  cm.3  of  50  per  cent  acetic  acid.  Dissolve  4  grm.  of  pure 

*  Tread  well-Hall:  Analytical  Chemistry,  3rd  ed.,  Vol.  II,  pp.  93-97. 
t  Gooch  and  Peters:  Methods  in  Analysis,  Gooch,  pp.  41,  42. 
j  Perkins:  Ibid.,  pp.  42,  361,  362. 

§  See  Methods  in  Analysis,  Gooch,  pp.  123,  131,  181-184,  187,  197,  218, 
248,  254. 


134 


QUANTITATIVE  CHEMICAL  ANALYSIS 


oxalic  acid  in  40  cm.3  of  50  per  cent  acetic  acid,  add  10  cm.3  of  concentrated 
nitric  acid,*  and  pour  this  solution  into  that  of  the  copper  sulphate.  Col- 
lect the  precipitate,  after  standing  over  night,  upon  asbestos  in  a  perforated 
crucible  and  wash  carefully  with  small  amounts  of  water  applied  succes- 
sively. Place  the  crucible  with  its  contents  in  a  beaker,  cover  it  with  about 
200  cm.3  of  hot  water  containing  25  cm.3  of  dilute  sulphuric  acid  (1  :  3),  and 
titrate  the  solution  with  N/10  permanganate.  From  the  oxygen  value 
of  the  permanganate  used  calculate  the  copper  in  the  copper  oxalate, 
CuC2O4,  titrated. 

State  the  results  in  the  following  form: 


CuS04  •  5  H20 
taken. 

Copper  in 
CuS04  •  5  H20. 

N/10  KMnO4 
used. 

"  Oxygen 
value  "  of 
KMnO4. 

Copper 
found. 

Error. 

Grm. 

Grm. 

Cm.3 

Grm. 

Grm. 

Grm. 

Determination  of  Iron  in  Ferrous  Sulphate.  f  —  When 
ferrous  sulphate  is  acted  upon  by  potassium  permanganate  in 

*  The  acetic  acid  is  used  to  increase  the  insolubility  of  the  copper 
oxalate  and  the  nitric  acid  is  added  to  induce  a  favorable  condition  for 
the  crystallization.  For  a  discussion  of  this  procedure,  and  for  its  applica- 
tion in  the  separation  of  copper  from  certain  other  elements,  see  Methods  in 
Analysis,  Gooch,  p.  123. 

f  Instead  of  potassium  permanganate,  standard  potassium  dichromate 
(Penny's  Method)  may  be  used  as  the  oxidizing  agent  in  the  determination 
of  iron  in  the  ferrous  condition,  the  reaction  taking  place  in  the  cold  solu- 
tion and  in  presence  of  either  hydrochloric  acid  or  sulphuric  acid: 
K2Cr2O7  +  6  FeSO4  +  8  H2SO4 

=  2  KHSO4  +  Cr2(SO4)3  +  3  Fe2(SO4)3  +  7  H2O 


K2C207  +  6  FeCl2  +  14  HC1  =  2  KC1  +  2  CrCl3  +  6  FeCla  +  7  H2O. 
The  end  of  the  reaction  is  not  indicated  by  any  characteristic  change  in 
color  in  the  solution,  but  is  found  by  noting  the  point  where  a  drop  of  the 
solution  taken  out  on  the  end  of  a  clean  stirring  rod  and  added  to  a  drop 
of  a  freshly  prepared  solution  of  potassium  ferricyanide  fails  to  give  with 
the  reagent  the  blue  which  shows  the  presence  of  a  ferrous  salt.  The 
dichromate  method  is  less  accurate  than  the  permanganate  method  of 
oxidation,  but  it  may  sometimes  be  used  with  advantage  when  solutions 
are  clouded  by  the  presence  of  insoluble  colored  material  in  suspension 
which  would  have  the  effect  of  obscuring  the  permanganate  end-color. 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS 


135 


presence  of  free  sulphuric  acid,  oxidation  takes  place  according 
to  the  reaction 

2  KMn04  +  10  FeS04  +  9  H2S04  =  2  KHS04  +  2  MnSO4 
+  5Fe2(S04)3  +  8H2O. 

The  relations  between  the  oxygen  derived  from  the  permanga- 
nate and  the  iron,  ferrous  oxide,  and  ferric  oxide  involved  are 
shown  in  the  condensed  expression 

2  FeO  +  0  =  Fe203, 

and  from  the  " oxygen  value"  of  the  permanganate  used  the 
amounts  of  the  other  substances  may  be  calculated. 

In  the  practical  application  of  this  reaction  to  the  determi- 
nation of  iron  in  a  ferrous  salt  (the  method  of  Margueritte),  the 
solution  diluted  so  that  the  concentration  shall  not  exceed 
1  :  1000  in  respect  to  iron,  is  treated  with  10  cm.3  of  sulphuric 
acid  (1  :  1)  for  every  100  cm.3  of  solution,  and  is  then  titrated  in 
the  cold  with  N/10  potassium  permanganate. 

Experimental  Process:  Iron  in  Ferrous  Ammonium  Sulphate.  —  Weigh  out 
carefully  about  0.7  grm.  of  ferrous  ammonium  sulphate  (NH4)2Fe(SO4)2  • 
6  EUO  (Mohr's  salt),  and  dissolve  it  in  cold  water.  Dilute  the  solution  to 
a  volume  of  200  cm.3,  add  20  cm.3  of  sulphuric  acid  (1  :  1),  and  run  in 
N/10  potassium  permanganate  to  coloration.  From  the  "oxygen  value" 
of  the  permanganate  calculate  the  weight  and  percentage  of  iron  in  the 
salt  taken.  State  the  details  of  the  determination  in  the  following  form: 


Oxygen 

(NH4)2 
FeS04  •  4  H2O 
taken. 

Theory  for  iron. 

KMnO4 
taken. 

value  per 
cm.3  of 
KMnO4 

Iron  found. 

Error  in  iron, 
found. 

used. 

Grm. 

Grm. 

Per  cent. 

Cm.3 

Grm. 

Grm. 

Per  cent. 

Grm. 

Per  cent. 

Determination  of  Hydrogen  Dioxide.  —  Hydrogen  dioxide 
in  sufficiently  dilute  solution  (0.001  per  cent)  and  in  presence 
of  sulphuric  acid  at  sufficient  concentration  (1  per  cent)  may 


136  QUANTITATIVE  CHEMICAL  ANALYSIS 

be  determined  by  direct  titration  with  N/10  permanganate,  by 
means  of  the  following  reaction: 

2  KMn04  +  4  H2SO4  +  5  H202  =  2  KHS04 

+  2  MnSO4  +  8  H2O  +  5  O2*. 

Determination  of  Nitrous  Acid  or  Nitrites.  —  Nitrous  acid 
is  oxidized  to  nitric  acid  according  to  the  equation: 

2  KMnO4  +  5  HN02  +  3  H2SO4  =  K2SO4 

+  2  MnSO4  +  3  H2O  +  5  HN03. 

In  the  analysis,  the  nitrous  acid  taken  in  solution  in  concen- 
trated sulphuric  acid  (on  account  of  the  volatility  of  nitrogen 
oxides  from  the  aqueous  solution)  is  measured  from  a  burette 
into  the  solution  of  a  known  amount  of  N/10  permanganate, 
diluted  to  about  400  cm.3,  and  warmed  to  40  degrees,  the  de- 
coloration of  the  solution  being  the  end-point.  Towarti  the 
end  the  bleaching  takes  place  tardily  and  the  nitrous  acid  or 
nitrite  must  be  added  slowly.  Nitrites  may  be  similarly  de- 
termined by  adding  the  neutral  or  alkaline  solution  to  a  known 
amount  of  N/10  permanganate,  suitably  diluted  and  acidified 
with  sulphuric  acid,  to  the  point  of  decoloration. 

Determination  of  Manganese  in  Manganous  Sulphate.  — 
When  potassium  permanganate  is  added  to  a  hot  and  neutral 
solution  of  manganous  sulphate  a  reaction  takes  place  (the 
Guyard  reaction)  in  which  the  low  oxide  MnO  of  the  manganous 
salt  is  oxidized  at  the  expense  of  the  high  oxide  Mn207  of  the 
permanganate  with  the  formation  of  intermediate  hydrated 
oxides.  For  example, 

Mn2O7  +  3  MnO  =  5  MnO2, 
Mn2O7  +  8  MnO  =  5  Mn2O3, 
Mn207  +  13  MnO  =  5  Mn3O4. 

These  higher  hydrated  oxides  appear  to  be  manganous  acid  or 
salts  of  manganous  acid,  H2MnO3.  When  the  conditions  are 

*  Should  it  happen  in  thus  testing  a  solution  for  hydrogen  dioxide  that 
the  first  drop  of  the  permanganate  gives  a  permanent  color  to  the  solution 
more  sulphuric  acid  should  be  added  before  concluding  that  no  hydrogen 
dioxide  is  present. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          137 

favorable  and  a  suitable  salt  of  another  element  (e.g.,  zinc) 
is  present,  a  manganite  of  this  element  is  precipitated  which, 
though  variable  in  composition,  contains  the  manganese  in  the 
quadrivalent  form.  In  the  presence  of  zinc  oxide,  a  weak 
base,  the  free  acid  is  taken  up  and  the  precipitation  of  the  man- 
ganese is  made  complete.  The  manganese  of  manganous  sul- 
phate may  therefore  be  determined  by  titration  of  the  solution, 
heated  to  boiling  and  containing  zinc  sulphate  with  zinc  oxide, 
until  the  supernatant  liquid  remains  pink.  The  course  of  the 
reaction  may  be  represented  typically  by  the  following  equa- 
tions : 

4  KMnO4  +  5  ZnSO4  +  6  MnS04  +  14  H2O 

=  4  KHSO4  +  7  H2SO4  +  5  ZnH204(MnO)2, 
H2SO4  +  ZnO  =  ZnS04  +  H20. 

In  this  process  (Volhard's  Method)  two  equivalents  of  the  per- 
manganate, KMn04,  supply  three  equivalents  of  oxygen,  and 
the  relation  of  the  available  oxygen  to  the  manganese  in  man- 
ganous form  is  shown  in  the  condensed  expression: 

MnO  +  6  =  MnO2. 

The  " oxygen  value  per  cm.3"  of  the  permanganate  solution 
when  used  in  this  process  would,  if  the  reaction  were  perfectly 
exact,  be  three-fifths  of  that  found  by  the  standardization  in 
the  presence  of  acid.  Practically,  it  is  more  exact  to  standardize 
the  permanganate  solution  against  a  known  weight  of  anhy- 
drous manganous  sulphate,  treated  exactly  as  in  the  process  of 
titration  by  which  the  unknown  manganous  salt  is  to  be  de- 
termined.* 

Dilution  and  the  presence  of  free  acid  tend  to  limit  this 
reaction  and  this  is  why  the  solution  remains  clear  for  a  con- 
siderable time  at  the  end-point  of  titrations  by  permanganate 
in  presence  of  a  suitable  amount  of  sulphuric  acid  (10  cm.3  of 
the  [1  :  1]  in  100  cm.3  of  solution),  and  why  such  solutions  may 
ultimately  become  brown  and  cloudy  upon  long  standing  after 
the  end-point  is  passed. 

*  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed.,  p.  612. 


138  QUANTITATIVE  CHEMICAL  ANALYSIS 

Determination  of  Formates.  —  The  determination  of  for- 
mates may  serve  as  an  example  of  oxidation  by  permanganate 
in  alkaline  solution.  In  cold  acid  solutions  the  action  is  very 
slow  and  from  hot  acid  solutions  formic  acid  may  escape  by 
volatilization.  An  excess  of  sodium  carbonate  is  therefore 
added,  the  solution  is  heated  to  induce  the  settling  of  the 
precipitate  to  be  formed,  and  the  permanganate  is  run  in  until 
the  clear  liquid  above  the  settled  precipitate  shows  the  end- 
color. 

2  KMn04  +  3  NaOCOH  =  K2C03  +  Na2CO3 

+  NaHCO2  +  CO2  +  Mn02. 

The  Effect  of  Hydrochloric  Acid  in  Permanganate  Ti- 
trations.  —  Although  hydrochloric  acid  and  potassium  per- 
manganate do  not  react  by  themselves  with  evolution  o£  chlo- 
rine at  the  concentrations  employed  in  processes  of  titration, 
chlorine  is  evolved,  a  portion  of  the  permanganate  is  wasted, 
and  the  indication  for  iron  is  too  high  when  hydrochloric  acid 
is  present  during  the  titration  of  a  ferrous  salt.*  The  same  is 
true  to  a  less  degree  in  respect  to  the  titration  of  oxalic  acid, 
and  other  oxidizable  substances,  in  the  presence  of  hydro- 
chloric acid.  It  has  been  found,  however,  that  the  introduction 
of  a  suitable  amount  of  a  manganous  salt  into  the  solution 
before  titration  will  not  only  prevent  the  secondary  action  by 
which  chlorine  is  evolved  but  will  very  markedly  hasten  a  reac- 
tion of  oxidation  when  it  is  slow  to  start. 

That  the  hastening  of  the  process  is  due  to  an  intervention 
of  the  Guyard  reaction,  described  above,  by  which  intermedi- 
ate oxides  of  manganese  (Mn02,  Mn2O3,  Mn3O4,  more  sensitive 
than  the  permanganate  to  the  action  of  a  reducing  agent)  are 
produced  in  the  action  of  permanganate  upon  a  manganous 
salt,  is  clearly  shown  in  the  titration  of  oxalic  acid  by  per- 
manganate. When  the  permanganate  is  added  drop  by  drop 
to  the  hot  solution  containing  oxalic  acid  and  sulphuric  acid 

*  This  effect  may,  however,  be  regulated  and  the  error  eliminated,  by 
making  the  titration  in  very  dilute  solution  and  calculating  the  results  by 
use  of  a  standard  obtained  by  matching  the  permanganate  against  a  known 
amount  of  ferrous  salt  under  similar  conditions. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          139 

the  solution  remains  colored  for  several  seconds  after  the  ad- 
dition of  each  of  the  first  few  drops,  during  which  process  a 
manganous  salt  is  formed,  and  then  the  rate  of  decolorization 
increases  gradually  and  rapidly.  If  a  suitable  amount  of  man- 
ganous sulphate  is  added  to  the  solution  of  oxalic  acid  and 
sulphuric  acid  before  titration  the  color  of  the  permanganate 
disappears  immediately  upon  the  addition  of  the  first  drop  and 
the  titration  may  be  pushed  rapidly,  even  in  the  cold  solution; 
but  if  too  much  manganous  salt  be  present  the  end  reaction  is 
marked  by  the  advent  of  a  brownish-red  precipitate  (due  to 
the  Guyard  reaction)  rather  than  the  clear  pink  of  the  soluble 
permanganate. 

The  explanation  of  the  protective  action  of  a  manganous 
salt  in  respect  to  the  evolution  of  chlorine  on  the  process  of 
titration  is  not  so  obvious,  although  it  would  seem  probable 
that  the  manganous  salt  is  instrumental  in  destroying  an  oxi- 
dizer  which  is  formed  in  this  reaction  and  is  sensitive  enough  to 
act  upon  hydrochloric  acid  as  well  as  upon  the  substance  to  be 
oxidized.  Neither  permanganic  acid  nor  any  of  the  persistent 
oxides  of  the  Guyard  reaction  will  evolve  chlorine  from  hydro- 
chloric acid  in  the  concentrations  employed  in  the  processes  of 
titration.  Manganic  acid,  H2Mn04,  the  oxide  of  which  (MnO3) 
stands  next  in  the  series  of  manganese  oxides  to  that  (Mn2C>7) 
of  permanganic  acid,  undergoes  spontaneous  decomposition 
when  it  is  formed  by  the  action  of  sulphuric  acid  upon  an  alkali 
manganate,  with  the  production  of  the  more  stable  manganous 
acid  and  permanganic  acid, 

3  H2Mn04  =  2  HMn04  +  H2Mn03  +  H20. 

Moreover,  chlorine  is  evolved  when  hydrochloric  acid  is  added 
to  the  solution  of  an  alkali  manganate.  In  order  to  explain  the 
evolution  of  chlorine  in  a  permanganate  titration  it  is  sufficient, 
therefore,  to  make  the  not  unnatural  assumption  that  in  the 
decomposition  of  the  permanganic  acid  any  one  of  the  known 
oxides  of  manganese  between  Mn2O7  and  MnO  may  be  formed 
momentarily  as  a  transition  step,  all  reducible  by  the  substance 
to  be  oxidized.  If  the  oxide  of  manganic  acid,  MnO3,  is  formed 
it  will  act  upon  hydrochloric  acid  as  well  as  upon  another 
oxidizable  substance,  but  if  the  amount  of  the  other  oxidizable 
substance  is  sufficiently  preponderant  the  action  upon  hydro- 
chloric acid  will  be  insignificant.  A  manganous  salt,  capable 
of  acting  with  manganic  acid,  as  well  as  with  permanganic 
acid,  to  form  the  persistent  oxides  of  the  Guyard  reaction,  is 
such  an  oxidizable  substance. 


140  QUANTITATIVE  CHEMICAL  ANALYSIS 

In  using  a  manganous  salt  for  the  purpose  of  regulating  the 
velocity  of  the  oxidizing  reaction  or  for  preventing  action  upon 
hydrochloric  acid,  it  is  important  to  adjust,  according  to  the 
conditions,  the  amounts  used,  in  order  that  these  desirable 
effects  may  be  secured  without  interference  with  the  end- 
reaction  by  the  persistence  of  the  Guyard  reactions. 

Experience  has  shown  that  in  titrating  oxalic  acid,  1  grm.  of 
the  manganous  salt  (MnSO4  •  5  H2O  or  MnCl2  •  4  H20)  should 
be  present  in  order  to  push  the  reaction  at  the  ordinary  tem- 
perature with  reasonable  speed,  and  that  this  amount  is  enough 
to  so  affect  the  conditions  of  equilibrium  that  titrations  in 
presence  of  hydrochloric  acid  (5  cm.3  to  15  cm.3  of  the  strong 
acid)  in  moderate  volumes  (100 'cm.3  to  500  cm.3)  may  be  con- 
ducted safely  and  with  reasonable  rapidity,  with  or  without 
sulphuric  acid.  ^ 

In  titrating  a  ferrous  salt,  which  acts  more  rapidly  than 
oxalic  acid  upon  the  oxides  produced  in  the  Guyard  reaction, 
a  considerable  amount  of  the  manganous  salt  may  be  present 
without  interfering  with  the  regularity  of  the  titration.  When 
much  of  the  iron  salt  is  present  hydrochloric  acid  may  bring 
out  the  yellow  color  of  ferric  chloride  which  is  detrimental  to 
the  delicacy  of  the  end-reaction,  but  this  difficulty  may  be 
overcome  by  the  addition  of  phosphoric  acid.*  In  a  volume  of 
500  cm.3,  which  may  contain  as  much  as  15  cm.3  of  strong 
hydrochloric  acid,  2  grm.  of  the  manganous  salt  (sulphate  or 
chloride)  will  be  sufficiently  protective  and  4  cm.3  of  phosphoric 
acid  of  sp.  gr.  1.7  with  the  usual  amount  of  sulphuric  acid  (10 
cm.3  of  the  (1:1)  acid  to  100  cm.3  of  solution)  will  serve  to  cor- 
rect the  color. 

Oxidations  by  the  Permanganate  Solution  after  a 
Preliminary  Reduction 

In  processes  of  analysis,  based  upon  the  oxidation  of  the  sub- 
stance to  be  determined  after  reduction  to  a  condition  from 
which  potassium  permanganate  will  reoxidize  definitely  the 

*  Reinhardt.  See  Analytical  Chemistry:  Treadwell-Hall,  3rd.  ed., 
Vol.  II,  p.  607. 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS  141 

product  of  reduction,  the  product  of  reduction  may  be  definite  and 
reasonably  stable  on  exposure  to  air;  the  reduction  product  may 
be  sensitive  to  the  action  of  air;  the  action  of  the  reducer  may  be 
excessive,  the  over-reduction  being  corrected  by  the  use  of  an  oxidizer 
of  limited  power  before  the  titration  with  the  standard  oxidizer. 

Reduction  to  Stable  Condition 

Determination  of  Iron  in  Ferric  Salts.  —  The  reduction  of 
ferric  salts  for  the  purpose  of  determining  the  iron  by  sub- 
sequent reoxidation  from  the  ferrous  condition  with  standard 
permanganate  may  be  effected  by  any  one  of  several  reducing 
agents,  of  which  the  one  most  commonly  used  is  zinc.  Other 
available  reducing  agents  are  hydrogen  sulphide,  sulphur  di- 
oxide, stannous  chloride,  and  titanous  sulphate. 

I.  Reduction  by  Zinc.  —  In  the  reduction  of  the  ferric  salt  by 
zinc  in  presence  of  free  acid  both  the  metal  and  nascent  hydro- 
gen may  take  part  in  the  reduction: 

Zn  +  2  FeCl3  =  ZnCl2  +  2  FeCl2, 

Zn  +  2  HC1  =  ZnCl2  +  2  H, 
2  FeCl3  +  2  H  =  2  HC1  +  2  FeCl2. 

When  a  ferric  salt  is  accompanied  by  a  salt  of  titanium  (as 
frequently  happens  in  the  analysis  of  titaniferous  iron  ores  and 
rocks)  the  latter  is  also  reduced  by  the  action  of  zinc  in  presence 
of  acid  and  left  in  condition  to  vitiate  the  determination  of 
iron,  but  by  treatment  of  the  solution  containing  the  products 
of  the  reduction  by  zinc  —  the  ferrous  salt  and  the  titanous 
salt  —  with  bismuth  oxide  only  the  titanous  salt  is  oxidized, 
3  Ti2(S04)3  +  Bi203  +  3  H2S04  =  6  Ti(SO4)3  +  2  Bi  +  3  H20. 

After  filtering  off  the  precipitated  bismuth,  with  the  excess  of 
bismuth  oxide,  the  ferrous  salt  may  be  titrated  in  the  usual 
manner.  To  prevent  the  error  which  may  be  caused  by  the 
presence  of  the  titanous  salt  it  is,  therefore,  only  necessary  to 
shake  the  solution  containing  the  ferrous  and  titanous  salts 
reduced  in  the  flask  (A)  or  in  the  reductor  (B),  with  a  little 
bismuth  oxide,  allow  the  mixture  to  stand  a  few  minutes,  and 
filter  quickly  with  the  aid  of  the  suction  pump,  and  titrate  as 


142  QUANTITATIVE  CHEMICAL  ANALYSIS 

usual  the  ferrous  salt  in  the  filtrate.*  There  is  little  danger  of 
reoxidation  of  the  ferrous  salt  when  the  cold  solution  free  from 
dissolved  oxygen  f  is  allowed  to  stand  for  a  short  time. 

II.  Reduction  by  Hydrogen  Sulphide.  —  For  the  determination 
of  iron  in  ferric  salts  by  the  reduction  and  subsequent  titration 
of  the  ferrous  salt  the  method  of  reducing  by  zinc  is  not  ap- 
plicable if  other  substances  similarly  reducible  and  pxidizable 
are  present.     In  such  cases,  notably  in  the  determination  of 
iron  in  presence  of  titanic  acid  (derived  from  titaniferous  ores) 
it  is  sometimes  customary  to  use  hydrogen  sulphide  as  the  re- 
ducer, the  reduction  taking  place  according  to  a  reaction  like 
the  following: 

Fe2(S04)3  +  H2S  =  2  FeS04  +  H2S04  +  S. 

In  this  process,  the  acidified  solution  of  the  ferric  salt  is  placed 
in  a  stoppered  flask  provided  with  a  tube  for  the  introduction 
of  hydrogen  sulphide  from  a  generator  and  another  tube  for 
the  exit  of  unused  gas.  The  contents  of  the  flask  are  heated  to 
boiling,  hydrogen  sulphide  is  passed  in  until  it  has  completely 
decolorized  the  ferric  salt  and  is  then  replaced  by  carbon 
dioxide  until  the  excess  of  the  hydrogen  sulphide  has  been  ex- 
pelled from  the  boiling  solution  and  the  escaping  gas  no  longer 
blackens  lead  acetate  paper.  The  solution  is  cooled  in  the 
atmosphere  of  carbon  dioxide  and  the  ferrous  salt  is  titrated 
with  the  usual  precaution  to  dilute  well  and  to  add  a  manga- 
nous  salt  if  hydrochloric  acid  is  present.!  The  precipitate  of 
sulphur  (if  free  from  sulphides)  does  not  reduce  permanganate 
in  the  cold,  dilute  solution. 

III.  Reduction  by  Sulphur  Dioxide.  —  A  ferric  salt  is  not 
completely  reduced  by  sulphur  dioxide  if  much  free  acid  is 
present  in  the  solution.     The  free  acid  is   therefore  nearly 
neutralized  by  sodium  carbonate,  sulphur  dioxide  is  bubbled 
through  the  boiling  solution,  contained  in  a  reduction  flask 
similar  to  that  used  in  the  process  of  reduction  by  hydrogen 
sulphide,  until  the  reduction  is  completed  and  is  then  replaced 
by   carbon  dioxide  until  the  gas  escaping  from  the  boiling 
solution  will  no  longer  decolorize  a  solution  containing  dilute 
sulphuric  acid  and  a  drop  or  two  of  N/10  permanganate.    The 

*  Gooch  and  Newton:   Methods  in  Analysis,  Gooch,  p.  500. 
t  Peters  and  Moody:  Methods  in  Analysis,  Gooch,  p.  371. 
J  See  p.  138. 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS  143 

ferrous  salt  is  titrated  after  cooling  the  solution  in  an  atmos- 
phere of  carbon  dioxide. 

Fe2(S04)3  +  S02  +  2  H20  =  2  FeS04  +  2  H2S04. 

IV.  Reduction  by  Stannous  Chloride*  —  Ferric  chloride  in  hot 
solution  and  in  presence  of  hydrochloric  acid  is  easily  reduced 
by  stannous  chloride.  The  stannous  chloride  is  added  cau- 
tiously, drop  by  drop,  until  the  solution  is  decolorized.  After 
thorough  cooling  of  the  solution  mercuric  chloride  is  added, 
and  after  a  short  interval,  the  titration  is  made  very  slowly  in 
the  diluted  solution  containing  a  manganous  salt.f 

2  FeCl3  +  SnCl2  =  2  FeCl2  +  SnCl4, 
SnCl2  +  2  HgCl2  =  SnCU  + 


V.  Reduction  by  Titanous  Sulphate.  —  When  a  solution  of 
titanous  sulphate  is  added  to  the  cold  solution  of  ferric  sulphate 
the  reduction  of  the  ferric  salt  is  immediate. 

Fe2(S04)3  +  Ti2(S04)3  =  2  FeS04  +  2  Ti(S04)2. 

The  excess  of  the  titanous  salt  is  immediately  oxidized  upon 
shaking  the  solution  with  a  little  bismuth  oxide 

3  Ti2(S04)3  +  Bi2O3  +  3  H2S04  =  3  Ti(S04)4  +  2  Bi  +  3  H20, 

and,  after  filtering  off  the  bismuth  and  the  excess  of  bismuth 
oxide,  and  diluting  the  solution,  the  ferrous  sulphate  may  be 
titrated  with  permanganate  in  the  usual  manner.  t  This 
method  §  is,  of  course,  directly  applicable  to  the  determination 
of  iron  in  ferric  salts  containing  a  salt  of  titanium.  A  solution 
of  titanous  sulphate  of  convenient  strength  may  be  made  by 
mixing  20  grm.  of  commercial  titanic  acid  with  three  times  its 
own  weight  of  a  mixture  of  sodium  and  potassium  carbonates 
and  fusing  it  in  a  platinum  crucible,  treating  the  melt  (after 
being  finely  ground)  in  a  platinum  dish  with  hot  concentrated 
sulphuric  acid,  cooling,  diluting  a  little,  filtering  through 
asbestos,  treating  with  zinc  until  reduction  is  accomplished, 
and,  while  zinc  is  still  left  in  the  flask,  filtering  the  solution 
quickly  through  a  perforated  platinum  cone  into  about  two 
liters  of  freshly  boiled  water  contained  in  a  small  reservoir  con- 
nected with  a  burette  and  a  hydrogen  generator. 

*  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed.,  Vol.  II,  p.  609. 

t  See  p.  138. 

|  See  p.  134. 

§  Newton:  Methods  of  Analysis,  Gooch,  p.  502. 


144  QUANTITATIVE  CHEMICAL  ANALYSIS 

Experimental  Processes:  Iron  in  Hematite.  — Extraction.  Weigh  out  into 
each  of  several  small  beakers  a  1-grm.  portion  of  the  finely  ground  iron  ore. 
To  each  portion,  add  10  cm.3  of  concentrated  hydrochloric  acid  (and,  if 
the  ore  contains  carbonaceous  matter,  a  pinch  of  potassium  chlorate), 
cover  the  beaker  with  a  watch-glass,  and  digest  at  a  temperature  just  be- 
low the  boiling  of  the  liquid  until  there  seems  to  be  no  further  solvent 
action.  Very  few  iron  ores  are  completely  decomposed  by  hydrochloric 
acid,  and,  unless  it  is  known  that  the  iron  of  the  ore  under  examination 
may  be  completely  dissolved  by  hydrochloric  acid,  the  insoluble  residue 
should  be  treated  further  to  extract  the  iron  in  the  insoluble  residue.  To 
this  end,  remove  the  watch-glass  (washing  any  spatters  upon  it  into  the 
beaker),  evaporate  the  solution,  and  desiccate  the  residue  to  render  the 
silica  insoluble.  Treat  the  residue  with  hydrochloric  acid  (5  cm.3),  dilute 
with  water  (40  cm.3),  settle,  decant  as  much  as  possible  of  the  clear  liquid 
and  filter  the  rest  through  a  small  filter,  washing  repeatedly  and  thoroughly, 
but  with  as  little  water  as  may  be.  Ignite  the  filter  and  siliceous  residue 
in  a  platinum  crucible,  add  sulphuric  acid  (about  1  cm.3)  and  hydro- 
fluoric acid  (about  2  cm.3)  and  evaporate  off  the  silicon  fluoride  and  the 
hydrofluoric  acid,  to  the  fuming  point  of  the  sulphuric  acid.  Cool  the 
liquid,  dilute  with  a  little  water,  and  if  no  dark  specks  of  undecomposed 
ore  remain  add  the  liquid  to  the  hydrochloric  acid  solution.  Should  any 
undecomposed  ore  still  remain,  filter  the  solution  into  the  flask,  through  a 
little  filter,  made  with  rough  edges  (by  tearing  the  edge  off  the  rim),  so 
that  it  will  fit  the  funnel  tightly  even  when  it  is  covered  with  liquid. 
Ignite  the  filter  and  residue,  fuse  with  acid  potassium  sulphate  until  the 
latter  is  quite  liquid  and  fuming  freely.  Cool,  dissolve-  the  residue  in  hot 
water  and  a  few  drops  of  hydrochloric  acid,  and  add  the  solution  to  the 
main  portion  of  the  ferric  solution. 

Reduction  by  Zinc.  —  A.  If  the  zinc  at  hand  is  not  iron-free,  as  is  ordi- 
narily the  case,  it  must  be  used  in  (roughly)  weighed  amounts  and  the  ex- 
cess of  it  entirely  dissolved  by  the  addition  of  sufficient  acid.  In  this 
case,  correction  must  also  be  made  for  the  iron  content  of  the  zinc,  found 
by  dissolving  a  weighed  portion  of  it  in  acid  by  itself  as  de- 
scribed below,  and  titrating  the  solution  with  permanganate. 
To  the  cold  solution  contained  in  a  small  flask,  of  50  cm.3  to  75 
cm.3  capacity,  add  3  grm.  of  granulated  zinc,  close  the  flask 
with  a  diminutive  funnel  (Fig.  26),  and  introduce  hydrochloric 
acid  (if  necessary)  until  a  brisk  evolution  of  hydrogen  takes 
2g.  place.  Keep  up  the  rapid  evolution  of  hydrogen  by  warming 
judiciously  and  adding  more  hydrochloric  acid  as  the  occasion 
may  require,  until  the  characteristic  yellow  color  of  ferric  chloride  has  dis- 
appeared and  does  not  return  upon  the  addition  of  a  few  drops  more  of 
the  acid.  The  use  of  hydrochloric  acid  in  the  reduction,  rather  than 
sulphuric  acid,  is  desirable  because  the  color  of  ferric  chloride  is  deeper 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS 


145 


than  that  of  ferric  sulphate  and  the  vanishing  of  it  a  better  indication  of 
complete  reduction.  Add  gradually  dilute  sulphuric  acid,  30  cm.3  of  the 
(1  :  2)  acid,  and  when  the  zinc  has  completely  disappeared  wash  down 
the  neck  of  the  flask,  and  the  funnel,  with  a  fine  jet  of  water  until  the  flask 
is  nearly  filled.  Stand  the  flask  in  cold  water,  transfer  the  cool  solution 
to  a  large  volume  of  cold  water 
(best  contained  in  a  white  dish) 
and  make  up  the  solution  to  a 
liter.  Add  in  solution  1  grm. 
of  manganous  sulphate  and 
titrate  the  ferrous  salt  in  the 
solution  in  the  usual  way. 

B.  If    iron-free   zinc   is   at 
hand  the  process  of  dissolving 
the    excess    of    zinc    may    be 
omitted.     In  this  case,  add  to 
the  solution  of  the  ferric  salt 
contained  in  a  small  flask  (50 
cm.3   to    75    cm.3)    a    suitable 
amount  of  the  iron-free  granu- 
lated zinc,  close  the  flask  with 
a    small    funnel,    add    hydro- 
chloric acid  (if  necessary)  until 
the  evolution  of  hydrogen  is 
brisk,  and  keep  up  the  action 
until  the  characteristic  color  of 
ferric  chloride  has  disappeared. 
Then,  while  the  evolution  of 
hydrogen   continues,   cool   the 
flask  hi  water,  strain  the  solu- 
tion  through   a  funnel   fitted 
with  a  platinum  cone  of  per- 
forated foil  or  gauze.     Add  the 
washings  of  the  flask  and  the 
zinc,  dilute  the  solution,  add 
manganous  sulphate  (1  grm.) 

and  syrupy  phosphoric  acid  (4  cm.3  of  sp.  gr.  1.7),  and  titrate  the  ferrous 
salt  in  the  usual  manner  by  permanganate. 

C.  When  many  reductions  are  to  be  made  the  work  is  very  much 
facilitated  by  the  use  of  the  Jones  redactor.     This  instrument,  shown  in 
Fig.  27,  is  so  arranged  that  the  reduction  may  be  effected  by  means  of  a 
column  of  granulated  zinc,  20  to  30  mesh,  —  that  is,  in  grains  which  will 
pass  the  sieve  of  twenty  meshes  to  the  linear  inch  or  eight  to  the  centimeter, 
but  will  not  pass  the  sieve  of  thirty  meshes  to  the  inch  or  twelve  to  the 


FIG.  27. 


146  QUANTITATIVE  CHEMICAL  ANALYSIS 

centimeter  —  previously  amalgamated  by  first  cleaning  it  with  a  little 
hydrochloric  acid  and  then  adding  mercuric  chloride  in  solution  until  the 
evolution  of  hydrogen  ceases.  This  amalgamated  zinc,  though  scarcely 
attacked  by  dilute  hydrochloric  or  sulphuric  acid,  is  capable  of  easily 
effecting  the  reduction  of  ferric  salts,  and  other  reducible  substances. 
The  reductor  tube  to  contain  the  zinc  has  (for  this  purpose)  a  length  of 
40  cm.,  and  an  inside  diameter  of  18  mm.  It  is  constricted  at  one  end  and 
sealed  to  a  stop-cock  tube.  The  tube  is  set  in  vertical  position,  a  cone  of 
perforated  platinum  foil  or  a  spiral  or  platinum  wire  is  inserted  (or  a  layer 
of  glass  beads)  upon  which  is  placed  a  plug  of  asbestos  or  glass  wool.  Above 
the  plug  is  put  a  thin  layer  of  asbestos  suitable  for  use  in  the  perforated 
crucible,  and  the  tube  is  filled  with  the  amalgamated  zinc  to  a  point  5  cm.3 
from  the  top.  The  reductor  tube  and  the  flask  to  receive  the  reduced 
solution  are  arranged  so  that  the  reduced  liquid  may  be  introduced  into 
water  without  contact  with  air,  as  shown,  and  connection  is  made  with 
the  suction  pump  and  a  pressure  regulator  which  will  keep  the  pressure 
inside  less  than  the  atmospheric  pressure  by  an  amount  equal  to  a  20-cm. 
column  of  water.  * 

In  making  use  of  the  Jones  reductor,  first  pour  warm  2.5  per  cent 
sulphuric  acid  into  the  tube  until  the  zinc  is  well  covered,  start  the  pump, 
adjust  the  stop-cock  so  that  the  liquid  will  run  slowly  into  the  receiver 
(best  cooled  in  running  water)  and,  keeping  the  zinc  always  covered  by  the 
liquid  (so  that  hydrogen  dioxide  may  not  be  produced  by  the  action  of  the 
nascent  hydrogen  upon  the  oxygen  of  the  air),  pass  100  cm.3  of  the  dilute 
acid  through  the  column.  Next  pass  the  solution  of  the  ferric  salt,  diluted 
with  100  cm.3  of  warm  2.5  per  cent  sulphuric  acid,  through  the  column  and 
wash  the  zinc  with  200  cm.3  of  the  warm  dilute  acid  followed  by  100  cm.3  of 
hot  water.  Cool  the  solution,  dilute  with  boiled  water,  and,  after  adding 
manganese  sulphate  (since  hydrochloric  acid  is  present),  titrate  in  the 
usual  manner  with  permanganate. 

The  presence  of  traces  of  iron  in  the  zinc  will,  of  course,  introduce  error 
into  determinations  made  with  the  Jones  reductor,  and  all  determinations 
made  with  this  instrument  should  be  corrected  by  taking  into  consider- 
ation the  reducing  effect  noted  in  a  blank  determination  in  which  the  other 
reagents  are  run  through  the  reductor  without  the  ferric  salt  and  titrated 
with  potassium  permanganate. 

From  the  number  of  cubic  centimeters  (ri)  of  the  permanganate  used 
and  its  oxygen  value  per  cm.3  (o),  the  weight  (w)  of  the  iron  indicated  may 
be  found  by  means  of  the  expressions: 


2FeO+ 

no  :  w  =  O:  2  Fe  =  16  :  111.68, 

no  X  111.68 
„=-_—-  • 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS 
Calculate  and  state  results  in  the  following  form: 


147 


Iron  ore  taken. 

KMnO4  used. 

Oxygen  value 
per  cm.3  of 
KMnO4 

Iron  found. 

A 

Grm. 

Cm." 

Grm. 

Grm. 

Per  cent. 

B 

C 

Determination  of  Vanadic  Acid  reduced  by  Sulphur  Di- 
oxide. —  When  sulphuric  acid  is  added  to  a  solution  which 
contains  vanadic  acid  or  a  vanadate,  the  passing  of  sulphur 
dioxide  into  the  boiling  solution  will  bring  about  the  reduction 
of  the  vanadic  acid  to  the  blue  vanadyl  salt,  corresponding  to 
vanadium  tetroxide,  which  is  perfectly  stable  in  the  air: 

V205  +  SO2  =  S03  +  V204. 

The  excess  of  sulphur  dioxide  is  removed  by  boiling  the  solu- 
tion, while  replacing  the  current  of  sulphur  dioxide  by  a  current 
'of  carbon  dioxide,  until  the  escaping  gas  no  longer  bleaches 
permanganate.  The  hot  solution  is  titrated  by  N/10  per- 
manganate and  the  calculation  is  made  according  to  the  equation 
V2O4  +  6  =  V2O5. 

When  small  amounts  of  vanadium  are  involved  the  pink  end- 
point  is  easily  recognizable  and  the  determinations  are  fairly 
accordant  and  exact.  Wider  variations,  due  to  the  difficulty 
of  catching  the  pink  end-reaction  in  presence  of  the  reddish 
yellow  color  of  vanadic  acid  at  high  concentration,  are  in- 
herent in  the  permanganate  process  of  titration  when  large 
amounts  of  vanadic  acid  are  formed. 

Reduction  Products  Sensitive  to  the  Action  of  Air 
In  certain  processes  the  reduction  is  definite  but  the  product 
is  so  sensitive  to  the  action  of  the  atmospheric  oxygen  that 
direct  titration  by  the  permanganate  solution  in  air  leads  to 
incorrect  results.  In  such  cases  the  reduced  solution  may  be 
received  directly  in  the  solution  of  a  reagent  which  will  im- 
mediately oxidize  it  with  the  simultaneous  formation  of  a 
product  which  is  less  sensitive  to  atmospheric  action,  though 


148  QUANTITATIVE  CHEMICAL  ANALYSIS 

easily  acted  upon  by  the  permanganate  solution  and  capable 
of  acting  as  a  register  of  the  reduction  previously  brought  about 
in  the  substance  to  be  determined.  Examples  of  such  processes 
are  the  determination  of  molybdic  acid  reduced  by  the  zinc 
reductor  and  received  in  ferric  sulphate,  and  the  similar  de- 
termination of  vanadic  acid  *  and  titanic  acid.f 

Determination  of  Molybdic  Acid  and  Phosphomolybdic 
Acid.  —  Molybdic  acid  may  be  reduced  by  zinc  in  the  presence 
of  sulphuric  acid  to  the  condition  of  the  oxide  IVfeOs  but  this 
condition  is  not  stable  in  presence  of  air.  When,  however,  the 
solution  containing  molybdic  acid  and  sulphuric  acid  is  passed 
through  the  Jones  reductor  {  and  then  directly  into  a  solution 
of  ferric  sulphate  placed  in  the  receiver,  the  sensitive  molybdic 
compound  is  partially  oxidized  by  the  ferric  salt  with  the 
formation  of  a  corresponding  amount  of  ferrous  salt,*  less 
sensitive  than  the  molybdic  compound  to  the  atmospheric 
actions.  §  The  molybdenum  compound  is  green  as  it  passes 
through  the  lower  part  of  the  reductor  but  on  coming  in  con- 
tact with  the  ferric  salt  it  is  changed  to  a  bright  red.  The 
solution,  decolorized  by  phosphoric  acid,||  is  titrated  while  still 
hot  with  N/10  permanganate.  The  result  may  be  calculated 
from  the  expression 

Mo203  +  30  =  3  Mo03. 

The  molybdenum  trioxide  contained  in  ammonium  phos- 
phomolybdate  may  be  similarly  reduced  and  determined  by 
the  procedure  first  outlined;  and  from  the  amount  of  per- 
manganate used  in  the  titration  of  ammonium  phosphomolyb- 
date  precipitated  under  definite  conditions  the  phosphoric 
pentoxide,  or  the  phosphorus,  contained  in  the  precipitate  may 
be  estimated. 

Experimental  Process:  Determination  of  Phosphorus  in  Iron  and  Steel.  — 
Of  steel  drillings  weigh  out  1  grm.  or  2  grm.  (according  as  the  steel  con- 
tains more  or  less  than  0.15  per  cent  of  phosphorus)  into  a  250-cm.3  Erlen- 

*  Gooch  and  Edgar:  Methods  in  Analysis,  Gooch,  p.  349. 

t  Newton:  Ibid.,  p.  242. 

t  See  p.  145. 

§  Randall:  Methods  in  Analysis,  Gooch,  p.  424. 

||  See  p.  140. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS 


149 


meyer  flask.  Add  a  mixture  of  25  cm.3  of  nitric  acid  (sp.  gr.  1.42)  and 
75  cm.3  of  water.  Hang  in  the  neck  of  the  flask  a  small  funnel  and  heat 
until  the  metal  has  dissolved  and  the  nitrogen  oxides  have  been  expelled. 
Add  in  concentrated  solution  0.15  grm.  of  potassium  permanganate  and 
boil  the  solution  until  the  characteristic  permanganate  color  has  disap- 
peared. Dissolve  the  precipitated  higher  oxides  of  manganese  by  adding 
a  solution  of  sulphurous  acid,  boil  out  the  excess  of  this  reagent,  and 
filter  the  solution.  Add  40  cm.3  of  ammonia  (sp.  gr.  0.96),  mix  thor- 
oughly to  dissolve  in  the  free  acid  the  precipitated  ferric  hydroxide,  and 
to  the  solution,  at  a  temperature  of  40  degrees,  add  40  cm.3  of  "molybdate 
solution,"  *  close  the  flask  with  a  rubber  stopper,  and  shake  vigorously  for 
five  minutes.  Settle  the  precipitate,  filter  (best  on  asbestos  in  a  perforated 
crucible),  and  wash  with  a  solution  of  ammonium  acid  sulphate  (15  cm.3 
of  ammonia,  25  cm.3  of  sulphuric  acid,  and  1  liter  of  water).  Wash  out 
the  flask  with  a  solution  of  20  cm.3  of  water  and  5  cm.3  of  ammonia,  and 
use  this  mixture  to  dissolve  the  precipitate  off  the  filter.  Acidify  the 
solution  and  washings  with  10  cm.3  of  strong  sulphuric  acid  and  pass  the 
acidified  solution  through  the  reductor  into  the  receiver,  containing  2  grm. 
to  3  grm.  of  ferric  alum  and  4  cm.3  of  syrupy  phosphoric  acid,  preceded  by 
100  cm.3  of  hot  water  and  followed  by  200  cm.3  of  the  hot  dilute  acid  with 
100  cm.3  of  water.  Titrate  the  reduced  solution  immediately  with  ap- 
proximately N/10  permanganate.  On  the  assumption  that  the  ammonium 
phosphomolybdate  contains  phosphorus  and  molybdenum  in  the  propor- 
tion given  by  the  symbol  (NH4)3PO4  •  12  MoO3,  and  that  the  reduction 
proceeds  to  the  condition  represented  by  the  symbol  Mo2O3,  the  weight  of 
phosphorus  may  be  calculated  from  the  number  of  cubic  centimeters  of 
permanganate  used  (ri)  and  the  oxygen  value  per  cm.3  (o)  of  that  reagent. 
no  :  w  =  18  O  :  P  =  288  :  31. 
no  X31 


w 


288 


State  the  results  in  the  following  form: 


Weight  of  steel 
taken. 

KMn04  used. 

Oxygen  value  per 
cm.'  of  KMnO4. 

Weight  of  P 
found. 

Per  cent  of  P 
in  steel. 

Grm. 

Cm.3 

Grm. 

Grm. 

*  Made  by  stirring  100  grm.  of  molybdenum  trioxide,  Mo03,  into 
80  cm.3  of  concentrated  ammonia,  adding  400  cm.3  of  water,  filtering, 
pouring  the  solution  into  a  mixture  of  400  cm.3  of  nitric  acid  (sp.  gr.  1.42) 
with  600  cm.3  of  water,  treating  with  0.05  grm.  of  microcosmic  salt, 
HNaNH4P04  •  4  H2O,  standing  twenty-four  hours,  and  filtering. 


150  QUANTITATIVE  CHEMICAL  ANALYSIS 

Over-reduction  and  Correction 

Excessive  and  irregular  effects  of  a  reducing  agent,  em- 
ployed to  prepare  a  substance  for  determination  by  the  per- 
manganate oxidation,  may  sometimes  necessitate  corrective 
treatment  before  the  titration  is  made.  The  action  of  zinc 
upon  vanadic  acid  and  uranyl  salts  will  serve  to  illustrate  this 
point. 

Determination  of  Vanadic  Acid  Reduced  by  Zinc.  —  The 
reduction  of  vanadic  acid  by  zinc  in  presence  of  sulphuric  acid, 
for  example,  proceeds  farther  than  the  condition  of  the  oxide 
V204  and  under  favorable  conditions  may  reach  that  of  V2O2. 
The  lavender  solution  thus  obtained  turns  blue  on  exposure  to 
air,  in  consequence  of  the  formation  of  a  blue  vanadyl  salt 
which  corresponds  to  the  tetroxide;  but  the  action  of  molecular 
oxygen  is  insufficient  to  bring  about  complete  oxidation  of  the 
lower  unstable  oxides  to  the  condition  of  the  stable  oxide 
V204  within  a  reasonable  time.  Ordinary  oxidizers  attack  the 
tetroxide  as  well  as  the  lower  oxides.  It  has  been  found,  how- 
ever, that  silver  oxide  and  silver  salts  are  sufficiently  active  to 
affect  the  lower  oxides  while  leaving  the  tetroxide  intact,  and 
silver  sulphate  appears  to  be  the  most  convenient  form  in  which 
to  use  the  silver  compound.* 

When  the  solution  of  vanadic  acid  is  subjected  in  presence  of 
sulphuric  acid,  as  in  the  reduction  of  a  ferric  salt,  to  the  action 
of  a  40-cm.  column  of  the  Jones  reductor  and  is  received  in  a 
flask  charged  with  a  saturated  solution  of  silver  sulphate  a 
muddy  deposition  of  finely  divided  silver  is  formed.  Upon  boil- 
ing the  mixture  the  silver  gathers  into  a  single  spongy  mass  and 
leaves  the  solution  so  clear  that,  were  it  not  that  spongy  silver 
is  susceptible  to  the  action  of  the  permanganate,  the  titration  of 
the  vanadium,  now  definitely  in  a  condition  corresponding  to 
V2C>4,  might  be  made  without  previous  filtration.  The  solution 
is  filtered  upon  asbestos  in  the  perforated  crucible,  diluted  to 
about  700  cm.3,  heated  to  the  boiling  point,  and  titrated  with 
N/10  permanganate.  The  result  is  calculated  from  the  ex- 
pression 

V204  +  6  =  V205. 
*  Gooch  and  Gilbert :  Methods  in  Analysis,  Gooch,  p.  346. 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS          151 

Determination  of  Uranium  in  Uranyl  Sulphate.  —  When 
uranyl  sulphate  in  dilute  sulphuric  acid  is  subjected  to  the 
action  of  the  zinc  column  of  the  Jones  reductor,  reduction  takes 
place  to  a  degree  dependent  upon  the  conditions.*  The  sul- 
phuric acid  should  have  during  the  action  the  concentration  of 
the  (1:6)  acid,  since  with  less  concentration  of  acid  the  re- 
duction is  delayed  while  a  greater  concentration  produces  a 
too  violent  evolution  of  hydrogen.  With  the  acid  of  the  con- 
centration named,  fifteen  minutes  or  more  should  be  allowed 
for  the  passage  of  the  uranium  solution  containing  as  much  as 
0.2  grm.  of  uranic  oxide  and  for  0.3  gram  half  an  hour  or  more 
should  be  allowed.  The  reduction  is  carried  under  this  con- 
dition beyond  the  uranous  stage,  represented  by  the  oxide 
UO2,  and  if  the  reduced  solution  is  received  and  titrated  in  an 
atmosphere  of  carbon  dioxide  the  indication  for  uranous  oxide 
is  too  high.  Brief  exposure  of  the  solution  to  air,  as  it  comes 
from  the  reductor,  is,  however,  sufficient  to  correct  the  over- 
reduction,  and  the  contents  of  the  receiving  flask,  which  after 
the  reduction  are  olive-green,  immediately  become  sea-green 
(the  characteristic  color  of  uranous  salts)  when  poured  out  into 
an  open  dish.  In  the  titration  of  the  hot  solution  of  uranous 
sulphate  with  permanganate,  the  solution  becomes  more  and 
more  yellowish  green  as  the  highest  condition  of  oxidation  is 
approached.  If  only  small  amounts  of  uranium  are  present 
the  addition  of  a  single  drop  of  N/10  permanganate  in  excess 
will  bring  out  the  pink  end-color,  but  with  large  amounts  of 
uranium  the  end-color  is  a  yellowish  pink. 

The  uranyl  sulphate  solution  amounting  to  100  cm.3  or  150 
cm.3,  and  containing  enough  sulphuric  acid  to  make  the  con- 
centration 1:6,  is  heated  nearly  to  the  boiling  point.  Pre- 
ceded by  a  few  cubic  centimeters  of  acid  of  the  same  strength, 
the  solution  is  drawn  by  gentle  suction  very  slowly  through  the 
40-cm.3  column  of  20-inch  amalgamated  zinc,  and  is  followed 
by  more  of  the  same  acid  and  then  by  250  cm.3  of  hot  water. 
The  contents  of  the  receiving  flask  are  poured  through  the  air 
into  an  open  porcelain  dish,  diluted  with  about  200  cm.3  of  hot 

*  Pulman:  Methods  in  Analysis,  Gooch,  p.  430. 


152  QUANTITATIVE  CHEMICAL  ANALYSIS 

water,  and  titrated  with  N/10  potassium  permanganate.  The 
reaction  takes  place  according  to  the  equations 

UO2SO4  +  Zn  +  2  H2SO4  =  ZnSO4  +  U(SO4)2  +  2  H20, 
2  KMn04  +  5  U(SO4)2  +  2  H2O  =  2  KHS04  +  2  MNSO4 
+  H2SO4  +  5  UO2SO4, 

and  the  result  is  calculated  from  the  expression 
U02  +  6  =  U03. 

Oxidations  by  Permanganate  of  a  Reagent  Used  in  Excess 

Some  processes  rest  upon  the  oxidizing  action  of  the  substance 
to  be  determined  upon  a  suitable  reducing  agent,  the  excess  of 
which  may  be  determined  by  titration  with  the  permanganate  solu- 
tion, the  reduced  product  of  the  substance  being  inactive  toward  the 
permanganate.  Examples  of  such  processes  are  the  determina- 
tions of  the  higher  oxides  of  lead  and  manganese  by  means  of 
their  reaction  upon  oxalic  acid  or  upon  ferrous  sulphate,  the 
excess  of  which  may  be  determined,  and  determinations  of 
chloric  acid  and  of  nitric  acid  by  means  of  the  similar  reaction 
upon  ferrous  sulphate  under  carefully  defined  conditions. 

Determination  of  Higher  Lead  Oxides.  —  Lead  dioxide,  in 
presence  of  nitric  acid,  is  capable  of  acting  upon  oxalic  acid 
with  the  formation  of  lead  nitrate  which  is  not  susceptible  to 
the  oxidizing  action  of  the  permanganate  in  acid  solution: 

Pb02  +  H2C204  +  2  HN03  =  Pb(N03)2  +  2  H20  +  2  C02. 

The  other  higher  oxides  of  lead,  Pb3O4  (red  lead)  and  Pb203, 
are  resolved  into  lead  nitrate  and  hydrated  lead  dioxide  by 
the  action  of  2  N  nitric  acid: 

Pb304  +  4  HN03  =  2  Pb(N03)2  +  H20  +  H2Pb03. 

These  characteristics  make  it  possible  to  determine  the  purity 
of  lead  dioxide,  or  the  condition  of  oxidation  of  any  of  the 
higher  oxides  of  lead,  by  acting  with  nitric  acid  (8  per  cent) 
upon  the  oxide,  treating  with  a  known  amount  of  standard 
oxalic  acid  in  excess  of  that  required  for  the  reaction,  and 
titrating  in  the  hot  solution  the  excess  of  the  oxalic  acid  by 
means  of  standard  permanganate. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  153 

If  the  standard  of  the  oxalic  acid  solution  is  stated  in  terms  of  the 
oxygen  required  to  oxidize  1  cm.3  of  it  (its  "oxygen  equivalent  per  cm.3") 
and  the  standard  of  the  permanganate  is  given  in  terms  of  the  oxygen 
available  in  1  cm.3  of  it  (its  "oxygen  value  per  cm.2"),  the  calculation  of 
the  amount  of  lead  dioxide  indicated  is  simple. 

If    o  =  the  oxygen  value  of  the  permanganate  per  cm.3, 
o'  =  the  oxygen  equivalent  of  the  oxalic  acid  per  cm.3, 
n  =  the  number  of  cm.3  of  permanganate  used, 
n'  =  the  number  of  cm.3  of  oxalic  acid  used, 
x  =  amount  of  lead  dioxide  indicated, 
then  O  :  PbO2  =  16  :  239.1  =  rio'  -no  :x, 

239.1  (n'of  -  no) 
T6~ 

Determination    of    Higher    Oxides    of    Manganese.*  —  I. 

Reduction  by  Ferrous  Sulphate.  Manganese  dioxide  and  other 
high  oxides  of  manganese  react  with  ferrous  sulphate  in  presence 
of  sulphuric  acid  to  form  ferric  sulphate  and  manganous  sul- 
phate : 

Mn02  +  FeSO4  +  2  H2S04  =  MnS04  +  Fe2(S04)3  +  2  H20. 
If  the  total  oxygen  value  of  the  amount  of  ferrous  sulphate 
used  is  known  (that  is,  the  amount  of  oxygen  required  to  con- 
vert the  ferrous  salt  to  the  ferric  condition),  and  the  oxygen 
equivalent  of  the  excess  remaining  after  the  reaction  be  de- 
termined, the  difference,  A,  will  measure  the  active  oxygen  of 
the  amount  of  manganese  dioxide  taken.     The  weight,  w,  of 
manganese  dioxide  indicated  will  be  given  by  the  proportion 
0  :  Mn02  =  16  :  86.93  =  A  :  w, 
A  X  86.93 

-ar 

II.  Reduction  by  Oxalic  Acid.  Manganese  dioxide  is  attacked 
by  oxalic  acid,  in  presence  of  sulphuric  acid,  with  the  formation 
of  manganous  sulphate : 

MnO2  +  H2C204  +  H2S04  =  MnS04  +  2  H20  +  2  C02. 
Other  high  manganese  oxides  are  reduced  similarly  and  if 
the  reduction  is  brought  about  by  a  known  amount  of  oxalic 
acid  (N/5)  taken  in  excess  of  the  amount  which  will  take 
part  in  the  reaction,  the  determination  of  the  excess  of  oxalic 
*  Treadwell-Hall:  Analytical  Chemistry,  3rd.  ed.,  Vol.  II,  p.  624. 


154  QUANTITATIVE  CHEMICAL  ANALYSIS 

acid  will  provide  the  data  necessary  for  the  calculation  of  the 
available  oxygen  of  the  high  manganese  oxide  and  of  the  equiv- 
alent amount  of  the  dioxide  Mn02. 

If    o  =  the  oxygen  value  of  the  permanganate  per  cm.3, 
o'  =  the  oxygen  equivalent  per  cm.3  of  the  oxalic  acid, 
n  =  the  number  of  cm.3  of  permanganate  used, 
nf  =  the  number  of  cm.3  of  oxalic  acid  used, 
then  n'o'  —no  =  available  oxygen  of  the  MnC>2. 

If  x  =  weight  of  MnC>2  indicated, 
then  O  :  MnO2  =  16  :  86.93  =  n'o'^-  no  :  x, 

_  86.93  (n'o'  -  no) 
16 

Determination  of  Chlorates.  —  The  determination  of  a 
chlorate  may  be  effected  by  taking  advantage  of  its  reaction 
with  ferrous  sulphate,  in  presence  of  sulphuric  acid  in  boiling 
solution.*  The  chlorate,  in  amounts  up  to  0.5  grm.,  is  treated 
with  a  known  amount  of  approximately  N/5  ferrous  sulphate 
of  known  oxygen  equivalent,  in  excess  of  the  amount  theo- 
retically necessary  to  register  the  oxidizing  power  of  the  chlo- 
rate, according  to  the  reaction, 

MC103  +  6  FeS04  +  H2S04  =  MCI  +  3  Fe2(S04)3  +  3  H20, 

with  15  cm.3  of  sulphuric  acid  (1  :  3),  in  a  flask  trapped  to  pre- 
vent mechanical  loss.  The  mixture  is  brought  to  the  boiling 
point,  cooled  to  the  room  temperature,  diluted  to  a  volume  of 
about  600  cm.3,  and  titrated  with  N/10  permanganate  after 
the  addition  of  2  grm.  to  3  grm.  of  manganous  chloride. 

From  the  total  oxygen  value,  P,  of  the  permanganate  used  the  weight 
of  chlorate  is  found  from  the  proportion 

3  O  :  MC1O3  :P:w, 
P  X  1  grm.  mol.  MC1O3 
3X16 

Determination  of  Nitrates.  —  Soluble  nitrates,  in  presence 
of  sulphuric  acid,  are  decomposed  by  ferrous  sulphate  in  the 
sense  of  the  equation 

2  MN03  +  6  FeS04  +  5  H2S04  =  2  MHS04  +  3  Fe2(S04)3 

+  4H20  +  2NO. 
*  Phelps:  Methods  in  Analysis,  Gooch,  p.  258. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          155 

The  reaction  may  be  utilized  for  the  determination  of  a  nitrate 
by  the  estimation  of  the  amount  of  ferrous  sulphate  taken  and 
the  amount  of  that  reagent  remaining  after  the  action  is  com- 
pleted, provided  means  are  adopted  to  remove  the  nitrogen 
dioxide  so  that  it  may  not  act  as  a  carrier  of  oxygen  from  the 
air  to  the  ferrous  sulphate  and  may  not  itself  be  acted  upon  by 
permanganate  in  the  process  of  titration. 

The  estimation  of  nitrates  may  be  accomplished  successfully 
by  the  following  procedure.* 

The  apparatus  which  is  used  consists  of  a  250-cm.3  boiling 
flask,  closed  with  a  rubber  stopper  carrying  a  separating  funnel 
of  50  cm.3  capacity  to  serve  as  an  inlet  tube  and  a  glass  tube 
of  0.8-cm.  bore  to  serve  as  an  outlet  tube.  The  inlet  tube  is. 
constricted  at  the  lower  end,  and  the  outlet  tube  is  enlarged 
just  above  the  stopper  to  a  small  bulb  (to  prevent  mechanical 
loss  during  the  boiling)  and  bent  twice  at  right  angles.  The 
flask  is  supported  on  a  wire  gauze  above  a  Bunsen  burner  and 
the  outlet  tube  dips  under  mercury  contained  in  a  test-tube. 

Of  the  nitrate  to  be  analyzed,  amounts  up  to  0.5  grm.  are  in- 
troduced with  water  into  the  flask,  the  stem  of  the  separating 
funnel  being  left  full  of  water,  the  outlet  tube  is  adjusted  to 
just  touch  the  surface  of  the  mercury  in  the  tube,  and  the  air 
is  expelled  from  the  flask  by  boiling  the  solution  to  small  volume. 
A  definite  amount  of  standardized  ferrous  sulphate  solution 
known  to  be  in  excess  of  that  required  to  break  up  the  nitrate 
is  introduced  into  the  separating  funnel.  The  outlet  tube  is 
plunged  a  centimeter  or  two  deep  into  the  mercury  (which  is 
readily  accomplished  by  changing  the  position  of  the  flask  on 
the  wire  gauze),  and  the  flame  is  withdrawn  until  the  dimi- 
nution of  pressure  due  to  cooling  is  made  evident  by  the  rise  of 
mercury  in  the  outlet  tube.  By  applying  and  withdrawing 
the  flame,  while  regulating  inflow  of  the  solution,  the  ferrous 
salt  may  be  introduced,  without  admitting  air.  The  funnel  is 
carefully  washed  with  a  volume  of  hydrochloric  acid  nearly 
enough  to  equal  that  of  the  liquid  previously  in  the  flask  and 
the  washings  are  similarly  introduced.  After  the  pressure  has 
been  restored  in  the  apparatus  by  heating  the  flask,  the  outlet 
tube  is  again  raised  to  the  surface  of  the  mercury,  and  the 
solution  in  the  flask  is  boiled  down,  beyond  the  point  at  which 
the  dark  compound  of  nitrogen  dioxide  and  ferrous  sulphate  is 

*  Phelps:  Methods  in  Analysis,  Gooch,  p.  462. 


156  QUANTITATIVE  CHEMICAL  ANALYSIS 

broken  up,  to  a  volume  of  10  cm.3  to  15  cm.3.  After  dilution 
with  600  cm.3  of  cool,  boiled  water  and  addition  of  2  grm.  to  3 
grm.  of  manganous  chloride,  the  ferrous  salt  remaining  is 
titrated  with  N/10  permanganate.  The  weight  of  nitrate,  w, 
may  be  calculated  from  the  total  oxygen  value,  P,  of  the  per- 
manganate used,  by  means  of  the  proportion 


_  P  X  2  grm.  mol.  MNO3 
3X16 

IODOMETRIC   PROCESSES 

The  measurement  of  the  iodine  involved  in  definite  reactions 
affords  a  basis  for  many  valuable  processes  of  volumetric 
analysis.  Iodine  may  be  used  in  standard  solution  to  bring  about 
a  definite  reaction;  or  iodine  may  be  a  product  of  reaction  and 
determined  by  some  suitable  standardized  reagent.  Analytical 
processes  which  depend  upon  the  measurement  of  definite 
amounts  of  iodine  used  as  a  reagent  or  obtained  as  a  product 
of  reaction  are  known  as  iodometric  processes. 

The  estimation  of  sulphurous  acid  by  titration  with  a  stand- 
ard solution  of  iodine  in  potassium  iodide  until  the  solution 
shows  color,  according  to  Dupasquier's  original  method, 

I2  +  H2S03  +  H2O  =  2  HI  +  H2S04, 

is  typical  of  reactions  of  the  first  sort.  The  estimation  of  free 
chlorine  by  the  use  of  a  standardized  solution  of  sulphurous 
acid  for  the  purpose  of  bleaching  the  iodine  set  free  when  the 
chlorine  comes  into  contact  with  potassium  iodide  may  serve 
as  an  example  of  the  second  sort: 

(a)  C12  +  2  KI  =  2  KC1  +  I2, 

(6)  H2S03  +  I2  +  H2O  =  H2S04  +  2  HI. 

Iodine  and  Sulphurous  Acid.  —  As  a  matter  of  fact  this 
illustrative  method  reaction  of  iodine  with  sulphurous  acid  is 
only  realized  ideally,  as  Bunsen  found,  when  the  iodine  is 
presented  to  sulphurous  acid  of  a  concentration  no  greater 
than  0.04  per  cent.  Nevertheless,  upon  the  use  of  standard 
iodine  and  standard  sulphur  dioxide  in  suitably  dilute  solution 
(standardized  by  comparison  with  the  standard  solution  of 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  157 

iodine)  Bunsen  founded  many  determinative  methods  resting 
upon  the  estimation  of  iodine  set  free  from  potassium  iodide 
either  by  the  direct  action  of  the  determined  substance  upon 
that  reagent  or  by  the  action  of  the  chlorine  liberated  in  a  re- 
action between  the  substance  and  hydrochloric  acid. 

In  Bunsen's  procedure  the  standard  iodine  was  made  by 
dissolving  a  known  weight  of  purified  and  dried  iodine  in  a  con- 
centrated solution  of  potassium  iodide,  or  by  dissolving  simi- 
larly a  known  weight  of  ordinary  iodine  of  value  determined 
by  analysis.  Dilute  sulphurous  acid,  of  about  0.03  per  cent 
strength,  was  standardized  against  a  solution  of  iodine  of 
known  value.  Potassium  iodide  was  made  up  for  use  in  a 
solution  of  10  per  cent  strength.  Dilute,  clear  starch  solution 
was  used  to  indicate  the  presence  of  an  excess  of  iodine  in  the 
process  of  titration.  With  these  reagents  Bunsen  demon- 
strated a  number  of  typical  processes,  such  as  the  determination 
of  free  sulphurous  acid  (of  strength  not  exceeding  0.04  per 
cent);  of  iodine  (dissolved  in  potassium  iodide);  of  chlorine 
and  bromine  (allowed  to  act  upon  potassium  iodide  to  set  free 
the  equivalent  amount  of  iodine);  of  chlorites  and  hypochlo- 
rites  (decomposed  by  hydrochloric  acid  and  potassium  iodide, 
with  evolution  of  iodine);  of  chromates,  chlorates,  and  the 
higher  oxides  of  lead,  manganese,  nickel,  and  cobalt  (by  treat- 
ing each  of  these  substances  in  a  diminutive  distillation  flask 
with  hydrochloric  acid,  collecting  the  evolved  chlorine  in  potas- 
sium iodide,  and  estimating  the  iodine  set  free).  Bunsen 
pointed  out  that  iodates,  vanadates,  selenates,  manganates, 
ferrates,  ozone,  etc.,  should  be  determinate  by  similar  pro- 
cedure. Cerium  was  determined  in  presence  of  lanthanum  by 
a  chlorine  oxidation  in  alkaline  solution,  treatment  of  the 
precipitate  containing  cerium  dioxide  with  hydrochloric  acid  in 
the  distillation  flask,  absorption  of  the  chlorine  evolved  in 
potassium  iodide,  and  titration  of  the  liberated  iodine.  Fer- 
rous salts  and  arsenious  acid,  or  arsenites,  were  oxidized  by 
known  amounts  of  potassium  chromate  and  the  excess  of  the 
last,  determined  by  the  procedure  indicated  for  chromates, 
gave  by  difference  the  amount  used  in  the  oxidation. 


158  QUANTITATIVE  CHEMICAL  ANALYSIS 

These  demonstrations  were  sufficient  to  point  the  way  to  the 
iodometric  estimation  of  many  substances  capable  of  acting 
upon  potassium  iodide  with  liberation  of  iodine  or  of  acting 
upon  hydrochloric  acid  with  evolution  of  chlorine  which  is  de- 
terminable  from  the  iodine  set  free  when  that  reagent  is  brought 
into  contact  with  potassium  iodide. 

The  chief  inconveniences  of  Bunsen's  procedure  —  the  neces- 
sity of  using  very  dilute  solutions  of  sulphurous  acid  in  the 
titration  of  iodine  and  the  tendency  of  the  standard  sulphurous 
acid  to  change  in  strength  on  account  of  spontaneous  and  con- 
tinuous oxidation  —  were  obviated  when  Schwarz  replaced  the 
sulphurous  acid  by  sodium  thiosulphate. 

Iodine  and  Thiosulphate.  —  Sodium  thiosulphate  acts 
definitely  with  iodine,  in  neutral  or  faintly  acid  solutions  of  any 
reasonable  concentrations,  with  the  formation  of  sodium  tetra- 
thionate  and  sodium  iodide: 

I2  +  2  Na2S203  =  Na2S406  +  2  Nal. 

Though  definite  and  complete  in  the  neutral  or  acid  solution, 
the  action  of  iodine  upon  sodium  thiosulphate  is  in  alkaline 
solution  irregular  and  excessive,  even  when  the  alkalinity  is 
due  to  a  bicarbonate. 

Iodine  and  Arsenic  Trioxide.  —  The  reaction  between 
iodine  and  arsenious  acid, 

I2  +  H3As03  +  H20  <=»  2  HI  +  H3As04, 

is,  however,  reversible  and  can  be  made  to  proceed  to  comple- 
tion, in  either  direction,  by  attention  to  the  conditions  of  action. 
If  the  hydriodic  acid  formed  in  the  reaction  is  neutralized  as 
fast  as  it  is  formed  the  forward  action  will  proceed  to  com- 
pletion, but  in  order  that  no  secondary  reaction  between  iodine 
and  the  neutralizing  reagent  may  intervene  the  latter  must  be 
properly  selected. 

When  iodine  and  an  alkali  hydroxide  meet  the  reaction  re- 
sults in  the  formation  of  an  iodide,  a  hypoiodite,  and  ultimate- 
ly an  iodate: 

61  +  6  KOH  =  3  KI  +  3  KOI  +  3  H2O 
=  5KI  +  KI03   +3H2O. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          159 

The  reaction  between  iodine  and  a  neutral  carbonate  is  similar, 
proceeding  to  an  equilibrium, 

21  +  2  K2C03  +  H2O  <=»  KI  +  KOI  +  KHCO3. 

But  upon  acid  carbonates  in  solution  iodine  has  no  appreciable 
action,  and  upon  a  mixture  of  monoalkali  arsenite  and  the 
dialkali  arsenite  —  which  exist  in  the  solution  saturated  with 
carbon  dioxide  —  the  only  action  of  iodine  is  that  of  simple 
oxidation. 

The  reaction  between  iodine  and  an  alkali  arsenite  (arsenic 
trioxide  in  presence  of  an  alkali  bicarbonate), 
I2+NaH2As03+  2  NaHCO3  =  NaH2AsO4+  2  Nal+  2  CO2+H2O, 

is  practically  as  complete  and  definite  as  that  between  iodine 
and  sodium  thiosulphate. 

The  ease  with  which  arsenic  trioxide  may  be  prepared  and 
weighed  in  a  state  of  purity,  the  stability  of  a  properly  made 
arsenite  solution  for  long  periods,  and  the  definiteness  of  the 
reaction  between  iodine  and  the  arsenite  solution  make  the  use 
of  pure  arsenic  trioxide  as  a  primary  standard,  in  place  of  pure 
iodine,  most  convenient  and  desirable.  But  the  relations  dis- 
cussed above  make  it  imperative  that,  in  making  up  the  stand- 
ard arsenite  solution,  care  be  taken  to  use  a  regulated  amount 
of  alkali  hydroxide  (or  to  neutralize  the  excess  subsequently) 
and  to  have  acid  carbonate  present  for  the  neutralization  of  the 
hydriodic  acid  formed  in  oxidation  of  the  arsenite  by  the 
iodine.*  By  comparison  with  the  arsenite  solution,  carefully 
made  and  used  as  primary  standard,  the  standard  of  iodine 
solutions  made  up  approximately  may  be  readily  fixed;  and,  by 
comparison  with  iodine  solutions  thus  standardized,  the  exact 
standard  of  solutions  of  sodium  thiosulphate,  made  up  approxi- 
mately, may  be  determined.  With  these  three  standard  solu- 
tions —  the  arsenite,  the  iodine,  and  the  thiosulphate  —  and  a 
properly  made  solution  of  starch  to  be  used  as  an  indicator  of 
free  iodine,  many  problems  of  quantitative  analytical  chemistry 
may  be  easily  and  accurately  solved. 

*  For  the  discussion  of  these  relations  from  the  point  of  view  of  the 
ionic  theory,  see  Washburn:  Jour.  Am.  Chem.  Soc.,  30,  21  (1908). 


160  QUANTITATIVE  CHEMICAL  ANALYSIS 

Iodine,  lodate,  and  Hydrochloric  Acid.  —  Iodine,  free  or 
evolved  in  reaction,  may  sometimes  be  conveniently  deter- 
mined by  conversion  to  colorless  iodine  chloride  formed  in  the 
reaction  of  this  iodine  with  suitable  concentrated  hydrochloric 
acid  and  an  oxidizer.  The  oxidizer  best  adapted  to  the  purpose 
is  potassium  iodate  in  standard  solution.* 

When  a  solution  of  potassium  iodide  and  potassium  iodate  is 
acidified  with  dilute  hydrochloric  acid  iodine  is  set  free,  ac- 
cording to  the  expression 

5KI  +  KI03  +  6HC1  =  6KC1  +  3  I2  +  3H2O. 

If,  however,  the  hydrochloric  acid  is  present  in  sufficient  con- 
centration (60  per  cent  of  the  concentrated  acid)  with  the 
proper  amount  of  iodate  iodine  enters  into  union  with  chlorine : 

KIO3  +  6  HC1  +  2 12  =  KC1  +  5 IC1  +  3  H20, 
the  entire  reaction  being  expressed  by  the  equation 

KI03  +  2  KI  +  6  HC1  =  3  KC1  +  3  IC1  +  3  H20. 

The  end-point  of  the  reaction  is  best  indicated  by  the  disap- 
pearance of  the  iodine  color  from  a  small  amount  of  chloroform 
(5  cm.3)  when  the  mixture  in  a  stoppered  bottle  is  thoroughly 
shaken  in  the  intervals  between  successive  additions  of  the 
standard  iodate. 

The  Starch  Indicator 

While  in  colorless  solutions  of  small  volume  (100  cm.3)  the 
yellow  color  produced  by  a  single  drop  of  N/10  iodine  is  dis- 
tinct, it  is,  nevertheless,  desirable  in  most  cases  to  increase  the 
delicacy  of  the  end-reaction  by  the  use  of  a  sensitive  indicator. 
The  "starch  blue"  which  is  formed  when  iodine  and  pure 
starch  come  together  in  presence  of  an  iodide  is  a  very  delicate 
indicator  for  free  iodine. 

The  End-Color.  —  In  the  literature  of  iodometric  titration 
frequent  mention  is  made  of  the  fact  that  a  red  color  may 
sometimes  appear,  as  well  as  the  blue,  when  starch  is  used  as 

*  Andrews:  Jour.  Am.  Chem.  Soc.,  25,  796  (1903). 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS          161 

an  indicator  for  iodine.  It  may  happen,  also,  that  in  certain 
processes  in  which  the  red  coloration  appears,  notably  the 
titration  of  an  arsenite  or  of  an  antimonite  (tartar  emetic),  in 
presence  of  an  alkali  bicarbonate,  loss  of  iodine  will  occur,  out- 
side the  main  reaction,  when  the  readings  are  delayed  beyond 
the  first  appearance  of  the  red  coloration.  It  has  been  found,* 
however,  that  neither  the  loss  of  iodine  nor  the  red  coloration 
takes  place  when  a  freshly  made  solution  of  pure  starch  is 
employed.  Whenever  it  is  impracticable  to  make  use  of  a 
fresh  and  pure  product  it  is  best  to  add  the  starch  indicator 
only  at  the  end-point  of  the  titration.  In  this  way  the  end- 
color  may  be  made  a  clear  blue,  even  with  an  impure  starch. 

It  appears  that  ordinary  starch,  or  a  once  pure  starch  which 
has  undergone  partial  hydrolysis,  is  liable  to  contain  amidulin, 
which  like  starch  gives  a  blue  coloration,  and  erythrodextrin 
which  gives  a  red  coloration  with  iodine;  these  substances  be- 
ing the  first  two  dextrins  produced,  in  the  order  mentioned, 
when  starch,  either  in  the  solid  state  -or  in  solution,  undergoes 
hydrolytic  change.  Amidulin  is  readily  changed  under  the  in- 
fluence of  an  oxidizer  to  erythrodextrin  and  to  this  action  must 
be  ascribed  the  simultaneous  production  of  the  red  coloration 
and  disappearance  of  iodine  sometimes  noted  in  titrimetric 
processes.  With  an  impure  starch,  the  reading  from  the  first 
permanent  color,  whether  red  or  blue,  is  nearest  to  the  cor- 
rect value;  and  the  titration  should  be  made  in  well-diluted 
solutions  —  150  cm.3  to  200  cm.3  —  since  the  production  of  red 
is  at  a  minimum  and  the  loss  of  iodine  small  at  high  dilutions. 

The  Influence  of  Iodides.  —  The  formation  of  the  starch 
blue,  as  well  as  the  delicacy  of  the  iodine  yellow,  is  conditioned 
by  the  presence  of  an  iodide,  and  the  influence  of  iodides  upon 
the  delicacy  of  the  end  reaction  is  obviously  a  matter  of  con- 
siderable importance  analytically.  Experience  shows  that  the 
presence  of  0.3  grm.  of  potassium  iodide  in  volumes  not  ex- 
ceeding 300  cm.3  is  sufficient  to  bring  out  the  maximum  sharp- 
ness of  readings  made  either  by  means  of  the  iodine  yellow  or 
by  the  starch  blue,  and  that  the  indications  are  then  quite  as 
delicate  at  the  ordinary  room  temperature  as  at  the  temper- 

*  Hale:  Methods  in  Analysis,  Gooch,  p.  29. 


162  QUANTITATIVE  CHEMICAL  ANALYSIS 

ature  of  ice-water.  The  addition  of  more  potassium  iodide 
(though  not  harmful,  within  reasonable  limits)  renders  the 
readings  no  sharper  for  any  given  volume  within  that  range  of 
dilution.  It  is  to  be  noted  that  the  delicacy  of  the  indication 
turns  also  upon  the  concentration  of  the  iodide  and  that  the 
amount  of  iodine  necessary  to  bring  out  the  starch  blue  in 
presence  of  a  suitable  amount  of  potassium  iodide  increases 
with  the  dilution. 

According  to  the  experimental  evidence*  it  would  appear 
that  the  group  KI  •  I4  is  characteristic  of  the  blue  starch  iodide 
formed  in  presence  of  potassium  iodide,  and  that  the  delicacy 
of  the  starch  blue  coloration  depends  upon  keeping  a  concen- 
tration of  potassium  iodide  suitable  for  the  production  of  this 
group.  On  the  other  hand,  the  excessive  concentration  of  the 
potassium  iodide  may  be  detrimental  since  it  tends  to  chajige 
the  group  KI  •  I4  into  two  groups  KI  •  I2,  characteristic  of  a 
starch  red  which,  however,  reverts  to  the  condition  of  the 
starch  blue  on  further  dilution  of  the  solution. 

Preparation  of  the  Starch  Solution.  —  The  statement  has 
been  made  that  preparations  of  starch  from  different  sources 
will  show  different  capacities  for  absorbing  iodine.  This  is  in- 
correct. Any  variety  of  pure  starch  (e.g.,  potato,  rice  or  arrow- 
root) will  serve  for  the  preparation  of  the  starch  indicator. 
The  " starch  solution"  may  be  made  by  grinding  a  suitable 
amount  of  starch  (2  grm.)  to  a  paste  with  a  little  cold  water, 
pouring  the  mixture  into  boiling  water  (1  liter),  and  con- 
tinuing the  boiling  for  five  or  ten  minutes  so  that  a  translucent 
solution  is  obtained.  Only  the  clear  supernatant  liquid,  or 
the  clear  solution  after  filtration,  is  to  be  used.  The  solution 
thus  prepared  will  soon  deteriorate  in  consequence  of  the 
growth  of  mould  and,  unless  fresh  solutions  are  to  be  prepared 
as  needed,  means  must  be  taken  to  keep  the  starch  in  con- 
dition for  use,  —  as  may  be  done  by  adding  10  mgrm.  of  mer- 
curic iodide  in  the  process  of  grinding  in  cold  water  (Gastine); 
or,  by  filling  into  small  bottles,  the  cold,  filtered  solution, 
sterilizing  by  heating  these  for  two  hours  in  a  water-bath  at  100 

*  Hale:   Methods  in  Analysis,  Gooch,  p.  37. 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS          163 

degrees,  closing  the  bottles  by  means  of  soft,  sterilized  stoppers, 
and  keeping  the  individual  bottles  to  be  opened  as  the  occasion 
may  require  (Stokes). 

The  Preparation  and  Use  of  Standards 

Arsenic  Trioxide,  a  Primary  Standard.  —  Carefully  sub- 
limed and  anhydrous  arsenic  trioxide  serves  as  the  most  exact 
primary  standard  in  iodometric  processes.  It  may  be  prepared 
by  subliming  the  commercially  pure  oxide  from  a  large  porce- 
lain crucible  (or  small  porcelain  dish)  upon  a  watch-glass,  reject- 
ing the  first  (or  successive)  portions  of  the  sublimate  in  case  it 
shows  the  red  color  of  selenium  or  the  yellowish  color  of  arsenic 
trisulphide.  Impure  arsenic  trioxide  may  also  be  purified  by 
dissolving  it  in  hot  hydrochloric  acid,  filtering  hot,  crystallizing 
out  the  pure  product,  drying  in  the  air,  and  resubliming. 

To  oxidize  a  gram-equivalent  of  arsenic  trioxide  (As203)  in 
alkaline  solution  four  gram-atoms  of  iodine  (I)  are  required  : 

H-  2  12  +  5  H20  =  2  H3O3AsO  +  4  HI. 


Solutions  of  the  same  degree  of  normality  will  be  obtained 
when  the  arsenic  trioxide  and  the  iodine  are  taken  in  the 


proportion  of       p3  :  I  =      y     :   126.92  =  4.948  :   12.692  for 
the  same  volume  of  solution. 

Preparation  of  N/10  Arsenite.  —  A  convenient  method  for 
the  preparation  of  the  N/10  solution  of  arsenite  is  to  dissolve 
4.948  grm.  of  pure  arsenic  trioxide  in  a  concentrated  solution  of 
4  grm.  of  potassium  hydroxide  and  make  up  the  solution  to  1 
liter  by  adding  100  cm.3  of  a  saturated  solution  of  potassium 
hydrogen  carbonate  and  enough  water  to  complete  the  volume. 
The  standard  of  the  arsenite  solution  is  most  conveniently  ex- 
pressed in  terms  of  the  iodine  which  is  equivalent  to  the  arsenic 
trioxide  contained  in  1  cm.3  of  the  solution.  The  "iodine 
equivalent  per  cm.3"  of  the  carefully  made  N/10  arsenite  solu- 
tion is  exactly  equal  to  the  "  iodine  value  per  cm.3"  of  the  N/10 
iodine  solution,  or  0.012692  grm.  The  arsenic  trioxide  dis- 
solves much  more  readily  in  the  alkali  hydroxide  than  in  the 
acid  carbonate  alone,  and  the  proportion  given  is  not  enough 


164  QUANTITATIVE  CHEMICAL  ANALYSIS 

to  convert  the  trioxide  entirely  into  dipotassium  hydrogen 
arsenite  but  more  than  enough  to  form  dihydrogen  potassium 
arsenite.  Should  more  alkali  hydroxide  be  used  in  effecting 
solution,  it  should  be  neutralized  by  a  suitable  amount  of  acid 
(sulphuric  acid  or  hydrochloric  acid)  before  the  final  addition 
of  the  acid  carbonate  in  excess;  but  it  is  important  to  obtain 
the  neutralization  without  the  addition  of  any  indicator  con- 
taining alcohol,  which  is  susceptible  to  the  action  of  iodine  in 
an  alkaline  solution.  The  neutralization  is  easily  accomplished 
without  the  use  of  any  indicator  if  attention  is  paid  to  the 
amount  of  alkali  hydroxide  employed.  With  this  standard 
arsenite  solution  the  standardization  of  iodine  dissolved  in 
potassium  iodide  is  effected  by  titration  with  or  without  starch 
as  an  indicator,  according  to  circumstances. 

The  Standard  Iodine  Solution.  —  To  prepare  a  solution  of 
iodine  a  weighed  amount  of  potassium  iodide  is  dissolved  in  as 
little  water  as  possible,  and  to  this  solution  is  added  a  weight 
of  iodine  not  greatly  exceeding  one-half  that  of  the  dissolved 
iodide.  Upon  shaking  the  mixture,  the  iodine  will  go  into  solu- 
tion, and  when  this  has  been  accomplished  the  solution  is  suit- 
ably diluted  with  water.  An  approximately  N/10  solution  of 
iodine  may  be  made  by  thus  dissolving  25  grm.  of  potassium 
iodide  in  a  liter  flask,  adding  12.7  grm.  of  iodine,  shaking  until 
the  iodine  is  completely  in  solution,  and  diluting  to  the  mark. 
The  solution  is  best  standardized  exactly  by  titration  against 
the  standard  solution  of  the  arsenite  in  presence  of  starch;  or 
it  may  be  standardized  against  a  thiosulphate  solution  of  pre- 
viously determined  standard.* 

Experimental  Process:  Standardization  of  Iodine  Against  N/10  Arsenite. 
—  Into  an  Erlenmeyer  beaker  containing  100  cm.3  of  water  draw  a  con- 
venient portion  (45  cm.3)  of  the  standard  N/10  arsenite  solution,  and  into 
this  run  from  a  burette  the  nearly  N/10  iodine  solution  to  be  standardized 
until  the  liquid  becomes  visibly  yellow  from  the  slight  excess  of  iodine. 
Then  add  2  cm.3  or  3  cm.3  of  starch  solution,  bleach  the  starch  blue  by  the 
addition  of  a  drop  or  two  of  the  standard  arsenite  and  then  restore  a  dis- 
tinct reading  color  by  the  careful  addition  of  the  iodine  solution,  the  sen- 
sitiveness of  the  color  reaction  being  greater  in  this  direction.  From  the 
''iodine  equivalent  per  cm.3"  (i  =  0.012692  grm.),  the  number  of  cubic 
centimeters  (n)  of  the  arsenite  solution,  and  the  number  of  cubic  centi- 
meters of  the  iodine  solution  diminished  by  the  amount  of  the  single  drop 

*  See  pp.  166,  167. 


PROCEDURES,  IN   VOLUMETRIC  ANALYSIS          165 

in  excess  (0.03  cm.3)  required  to  bring  out  the  starch  color  (n')>  the  "iodine 
value  per  cm.3"  of  the  latter  (i'}  may  be  found  from  the  expression 


.,  _  n  (0.012692) 

~^~ 

Iodine  solutions  of  lower  concentrations  —  for  example, 
N/100  solutions  —  may  be  easily  made  by  suitably  diluting  the 
standardized  iodine  solution  of  nearly  N/10  concentration. 

The  Standard  Thiosulphate  Solution.  —  Sodium  thiosulphate 
reacts  with  iodine  in  neutral  solution,  or  when  added  with 
stirring  to  the  slightly  acid  solution,  according  to  the  equation 

Na-0-S:;° 

9  T      T4_oNa-0\s  =  O  1 

^  "~        Na  —  S/b-O      —  fe 

I    +2Na-I 

S 


In  alkaline  solution  the  reaction  between  sodium  thiosulphate 
and  iodine  goes  farther  and  is  irregular. 

Sodium  thiosulphate  is  acted  upon  by  free  acids  with  the 
ultimate  formation  of  sulphurous  acid  and  sulphur: 

Na2S2O3  +  2  HX  =  2  NaX  +  H2S203 


It  is  even  susceptible  to  the  action  of  carbonic  acid,  and  in  a 
freshly  made  solution  is  subject  to  slight  decomposition  in  con- 
sequence of  the  action  of  dissolved  carbon  dioxide;  but  when, 
after  an  interval  of  some  days,  the  carbon  dioxide  has  become 
fixed  in  the  form  of  an  acid  carbonate  the  solution  will  keep  its 
value  for  a  considerable  time.  It  is  to  be  noted  that  if  any  sul- 
phurous acid  is  produced  by  the  decomposition  of  the  thiosul- 
phate this  acid  has  a  capacity  to  react  with  iodine  twice  greater 
than  that  of  the  thiosulphate  from  which  it  is  derived.  The 
amount  of  iodine  which  may  enter  into  the  reaction  with  the 


166  QUANTITATIVE  CHEMICAL  ANALYSIS 

thiosulphate  is  dependent  to  a  very  marked  degree,  therefore, 
upon  the  concentrations  of  the  thiosulphate  and  the  free  acid, 
as  well  as  upon  the  duration  of  the  reaction  and  the  tempera- 
ture. When  considerable  amounts  of  thiosulphate  are  to  be 
titrated  by  iodine  in  an  acidified  solution  the  conditions  under 
which  the  reaction  takes  place  must  be  carefully  guarded.  The 
concentration  of  the  acid  should  be  restricted,  the  temperature 
should  be  reduced  as  nearly  as  possible  to  0  degrees,  and  the 
iodine  should  be  added  promptly.  So  long  as  the  thiosulphate 
present  does  not  exceed  a  concentration  equivalent  to  that  of 
the  N/100  solution  a  rapid  titration  proceeds  in  the  cold  solu- 
tion with  fair  regularity  in  presence  of  hydrochloric  acid  in  con- 
centrations not  greater  than  that  of  the  N/4  solution  of  that 
acid.  In  most  iodometric  processes  involving  the  use  of  the 
thiosulphate  in  acid  solution  it  is  possible  to  add  that  reagent 
gradually  to  the  solution  so  that  it  may  be  acted  upon  by  io- 
dine in  the  normal  way  as  fast  as  it  is  introduced.  Under  such 
conditions  the  danger  of  interaction  with  acid  is  avoided  and  the 
titration  may  be  made  at  the  ordinary  labor atory'temperature. 
An  approximately  N/10  solution  is  made  by  dissolving  in  a 
liter  of  water  25  grm.  of  the  crystallized  sodium  thiosulphate, 
Na2S203  •  5  H2O,  about  one-tenth  of  the  gram-equivalent;  but 
the  standard  of  the  thiosulphate  may  not  be  permanent  until 
the  solution  has  been  allowed  to  stand  for  some  days.  The 
solution  may  be  standardized  by  titration  against  (I)  iodine  in 
a  solution  previously  standardized  against  the  standard  arse- 
nite,  or  (II)  against  iodine  specially  prepared  and  weighed  out, 
or  (III)  against  iodine  set  free  from  an  iodide  by  standard 
potassium  permanganate,  and  several  other  less  exact  methods 
have  been  proposed. 

Experimental  Process:  Standardization  of  Thiosulphate  against  Standard 
Iodine.  —  Into  an  Erlenmeyer  flask  containing  100  cm.3  of  water  draw  a 
convenient  portion  (45  cm.3)  of  the  standardized  solution  of  iodine  and 
run  in  the  thiosulphate  solution  to  be  standardized  until  only  a  pale  yel- 
low tinge  remains  of  the  original  iodine  color.  Add  starch  solution  (2  cm.3 
or  3  cm.3)  and  bleach  the  starch  blue  to  the  vanishing  point  by  the  careful 
addition  of  more  of  the  thiosulphate.  Find,  from  the  number  of  cubic 
centimeters  (n)  of  the  iodine  solution  of  known  "iodine  value  per  cm.3"  (i) 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  167 

and  the  number  of  cubic  centimeters  (n'~)  required  to  bleach  the  iodine  color 
with  starch,  the  "iodine  equivalent  per  cm.3"  (ir)  of  the  thiosulphate: 


Experimental  Process:  Standardization  against  Weighed  Iodine.  —  Pure 
iodine  may  be  prepared  from  commercial  iodine  (which  may  contain 
chlorine,  bromine,  and  cyanogen,  as  impurities)  by  mixing  it  with  potas- 
sium iodide  and  subliming  it  from  the  iodide  (which  will  remove  the  im- 
purities) and  then  resubliming  it  to  get  rid  of  traces  of  iodide  which  may 
be  carried  with  the  iodine  in  the  first  sublimation.  In  carrying  out  this 
operation,  heat  the  mixture  in  a  beaker  which  is  closed  by  a  conical  or 
cylindrical  flask  filled  with  luke-warm  water.*  Loosen  the  sublimed  iodine 
which  collects  upon  the  condenser  in  the  form  of  a  crystalline  crust  by  re- 
placing the  water  in  the  condenser  by  cold  water,  thus  chilling  and  con- 
tracting the  glass.  Remove  the  crust  of  iodine,  and  resublime  it  similarly 
by  itself  to  get  rid  of  traces  of  iodide.  Grind  it  in  an  agate  mortar  and  dry 
in  a  desiccator  (with  joints  ungreased)  over  calcium  chloride.  This  drying 
agent  is  preferable  to  sulphuric  acid  because  the  latter  absorbs  the  vapor 
of  iodine  and  may  then  evolve  contaminating  volatile  products.  The  des- 
iccator is  used  without  grease  lest  the  latter  form  with  the  iodine  vapor 
hydriodic  acid  to  be  absorbed  by  the  solid  iodine. 

In  weighing  out  the  iodine  for  use  as  the  standard  of  reference,  with 
necessary  precaution  against  volatilization,  place  a  portion  (0.4  grm.  to 
0.5  grm.)  in  a  weighed  tube,  provided  with  a  glass  stopper  and  containing 
potassium  iodide  (2  grm.  to  2.5  grm.)  moistened  with  a  very  little  water 
(0.5  cm.3),  put  the  stopper  in  place  and  weigh  again.  Loosen  the  stopper  of 
the  tube  containing  the  iodine  dissolved  in  the  concentrated  solution  of  the 
iodide  and  allow  the  tube,  followed  by  the  stopper,  to  slide  down  the  side 
of  an  Erlenmeyer  beaker  held  at  an  inclination  and  containing  200  cm.3  of 
water.  Titrate  the  solution,  thus  prepared,  with  precaution  against  vola- 
tilization of  the  weighed  iodine  with  the  thiosulphate  to  be  determined  and 
find^the  "iodine  equivalent  per  cm.3"  of  the  thiosulphate. 

Experimental  Process:  Standardization  against  Iodine  Set  Free  by  Stand- 
ard Permanganate.  —  When  potassium  permanganate  acts  upon  an  acidified 
excess  of  potassium  iodide,  iodine  is  set  free  according  to  the  equation 

2  KMn04  +  10  KI  +  16  HC1  =  12  KC1  +  2  MnCl2  +  8  H2O  +  5  12, 

and  the  relation  of  the  available  oxygen  to  the  iodine  evolved  is  shown  in 
the  expression 

5  O  +  10  HI  =  5  H2O  +  10  1, 

*  For  a  specially  designed  apparatus  see  Treadwell-Hall  :  Analytical 
Chemistry,  3rd.  ed.,  Vol.  II,  p.  646. 


168  QUANTITATIVE  CHEMICAL  ANALYSIS 

from  which  it  appears  that  the  "iodine  equivalent  per  cm.3"  (?>)  of  the 
permanganate  will  be  to  the  oxygen  value  per  cm.3  (o)  in  the  relation 
126.92  :  8. 

To  standardize  an  approximately  N/10  solution  of  sodium  thiosulphate 
by  the  use  of  standard  permanganate,  dissolve  a  suitable  amount  of  potas- 
sium iodide  (3  grm.)  in  water  (200  cm.3),  add  hydrochloric  acid  (5  cm.3), 
and  (if  the  solution  remains  colorless)  a  measured  amount  of  standard 
permanganate  (45  cm.3  N/10),  and  titrate  with  the  thiosulphate  solution 
the  iodine  set  free,  using  starch  as  the  indicator.  Should  the  addition  of 
acid  to  the  solution  of  the  iodide  set  iodine  free  before  the  permanganate  is 
added,  the  solution  may  be  made  colorless  again  by  the  cautious  addition 
of  the  exact  amount  of  the  thiosulphate  required  to  bleach  the  iodine. 
Thereafter  the  permanganate  may  be  immediately  added.  Of  course,  only 
the  amount  of  the  thiosulphate  added  after  the  permanganate  is  to  be 
taken  into  account  in  the  calculation  of  the  relation  between  the  perman- 
ganate and  the  thiosulphate.  From  the  number  of  cubic  centimeters  of 
the  permanganate  taken  (n),  the  number  of  cubic  centimeters  of  the  thio- 
sulphate (n'}  used  in  taking  up  the  iodine  liberated  by  the  permanganate, 
and  the  "oxygen  value  per  cm.3"  of  the  permanganate  (o),  the  "iodine 
equivalent  per  cm.3"  of  the  thiosulphate  (if),  may  be  calculated 

.,      no  (126.92  --  8) 

1     —  ; • 


The  Standard  lodate  Solution.  —  The  standard  solution  of 
iodate  may  be  made  up  by  dissolving  in  water  a  carefully 
weighed  amount  (conveniently  10.7  grm.  (about  ^V  mol.)  to 
the  liter)  of  pure  potassium  iodate,  KI03,  neutral  to  litmus. 
If  not  at  hand,  pure  potassium  iodate  may  be  made  *  by  dis- 
solving potassium  permanganate  (40  grm.)  in  hot  water  (1 
liter),  adding  potassium  iodide  (20  grm.),  heating  for  half  an 
hour,  adding  alcohol  drop  by  drop  until  decoloration  of  the 
excess  of  permanganate  takes  place,  filtering,  washing  the 
residue  with  hot  water,  adding  acetic  acid  to  plainly  acid  reac- 
tion, evaporating  to  small  volume  (50  cm.3),  cooling  to  solution, 
and  washing  the  crystallized  product  repeatedly  with  strong, 
pure  alcohol.  The  oxidation  of  the  iodide  takes  place  according 
to  the  equation 

KI  +  2  KMn04  +  H20  =  KI03  +  2  KOH  +  2  Mn02. 
*  Groger:  Zeit.  Anorg.  Chem.,  1994,  13. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          169 

In  determining  free  iodine  in  presence  of  strong  hydrochloric 
acid,  by  means  of  the  iodate  solution, 

KIO3  +  6  HC1  +  2 12  =  KC1  +  5 IC1  +  3  H20, 

the  relation  between  the  iodate  used  and  the  iodine  converted 
to  iodine  chloride  is  given  by  the  expression 

KI03  <>4I. 

When  the  action  takes  place  between  the  iodate  and  an  iodide 
in  presence  of  strong  hydrochloric  acid, 

KI03  +  2KI  +  6HC1  =  3KC1  +  3  IC1  +  3H2O, 

the  relation  between  the  iodate  and  the  iodine  combined  in  an 
iodide  is  given  by  the  expression 

KIO3  o  2 1. 

The  standard  of  an  iodate  solution  may  also  be  fixed  by  de- 
termining, as  in  the  process  of  standardizing  iodine  (see  p.  164), 
the  amount  of  iodine  which  is  set  free  when  a  given  amount  of 
the  iodate  is  acted  upon  in  slightly  acidified  solution  by  an  ex- 
cess of  potassium  iodide, 

KI03  +  5KI  +  6HC1  =  6KC1  -f  3H20  +  3  I2, 

in  which  reaction 

6  I  o  KI03. 

The  Determination  of  Iodine 

Iodine  in  free  condition  or  evolved  in  reaction  may  be  dissolved 
in  potassium  iodide  and  determined,  under  suitable  conditions, 
by  means  of  standard  arsenite  (as  in  the  experimental  standard- 
ization of  iodine  against  N/W  arsenite)  or  standard  thiosulphate 
(by  a  process  the  reverse  of  the  experimental  standardization 
of  thiosulphate  against  weighed  iodine) . 

'From  the  iodine  equivalent  per  cm.3  (i)  of  the  standard  solution  and 
the  number  of  cubic  centimeters  of  it  (n),  the  weight  (w~)  of  the  iodine  to 
be  determined  may  be  found  from  the  expression  w  =  ni]  or,  in  case  the 
iodine  is  determined  by  an  exactly  N/10  standard,  by  the  expression 
w  =  n  X  0.012692. 


170  QUANTITATIVE  CHEMICAL  ANALYSIS 

The  standard  iodate  solution  may  be  used  for  the  determina- 
tion of  free  iodine  in  a  strong  solution  of  hydrochloric  acid. 
See  pp.  178,  180. 

Processes  Based  upon  Oxidation  by  Standard  Iodine 

As  implied  in  the  standardization  of  iodine  against  arsenic 
trioxide  and  in  the  similar  standardization  of  sodium  thiosul- 
phate  against  iodine  determinations  of  arsenic  trioxide-  and  of 
thiosulphates  may  be  made  by  direct  titration  with  standard 
iodine. 

Direct  Titration  by  Iodine 

Determination  of  Arsenic  Trioxide.  —  The  determination 
'of  arsenic  trioxide  combined  in  the  form  of  an  alkali  arsenite 
is  made  by  titration  with  standard  iodine  in  solution  alkaline 
with  an  alkali  hydrogen  carbonate  precisely  as  in  the  determi- 
nation of  the  relation  between  iodine  and  the  standard  arsenite 
in  the  process  of  standardization  of  the  former.*  The  reaction 
takes  place  according  to  the  equation 

KH2AsO3  +  I2+  2  NaHC03  =  KHAsO4+  2  Nal  +  H2O  +  2  C02, 

and  the  relation  between  arsenic  trioxide  and  iodine  is  given 
by  the  expression 

As203  +  2 12  +  2  H20  =  As2O5  +  4  HI. 

The  solution  of  pure  starch  will  give  a  sharp  end-reaction,  while 
with  impure  starch  there  may  be  an  over  use  of  iodine  ac- 
companying the  production  of  reddish  coloration.  If  only  an 
impure  starch  is  available  the  reading  should  be  first  made 
without  starch  in  presence  of  enough  potassium  iodide  to 
bring  out  sharply  the  yellow  iodine  color  (see  p.  161).  The 
reading  with  starch  may  then  be  confirmed  by  adding  the 
starch  solution  which,  though  impure,  will  then  give  only  the 
pure  blue  color. 

Determination  of  Thiosulphates.  —  Soluble  thiosulphates 
may  be  determined  by  titration  with  iodine  in  the  neutral  or 
acid  (not  alkaline)  solution,  as  in  the  standardization  of  the 

*  See  p.  164. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          171 

sodium  thiosulphate  solution.*  When  occurring  in  association 
with  sulphides  or  sulphydrates  the  thiosulphate  may  be  titrated 
after  treatment  of  the  mixture  with  freshly  precipitated  cad- 
mium carbonate  and  nitration  of  the  solution  from  the  precipi- 
tated cadmium  sulphide  and  the  excess  of  cadmium  carbonate. 
Determination  of  Antimony  Trioxide.  —  Another  example 
of  oxidation  processes,  based  upon  a  direct  titration  by  standard 
iodine,  is  the  determination  of  antimony  in  the  trivalent  con- 
dition. In  determining  antimony  trioxide,  in  a  soluble  com- 
pound, the  titration  is  made  by  standard  iodine  in  presence  of 
an  alkali  hydrogen  carbonate  exactly  as  in  the  titration  of  an 
arsenite.  It  is  necessary,  however,  to  have  the  antimony  in 
such  combination  that  it  may  resist  precipitation  as  antimoni- 
ous  acid  (or  an  oxy-salt),  due  to  the  hydrolytic  action  of  water 
or  to  the  action  of  the  alkali  hydrogen  carbonate,  and  this 
purpose  is  accomplished  by  adding  tartaric  acid  or  an  alkali 
tartrate  (such  as  Rochelle  salt)  to  the  liquid  in  which  the 
antimony  compound  (e.g.,  antimony  trichloride)  is  to  be  dis- 
solved : 

Sb2O3  +  2  I2  +  2  H2O  =  Sb2O5  +  4  HI. 

Reverse  Titration 

Determination  of  Sulphurous  Acid.  —  According  to  Bunsen, 
when  the  concentration  of  the  sulphur  dioxide  is  very  low  — 
not  exceeding  the  limit  of  0.04  per  cent  of  the  solution  —  Du- 
pasquier's  method  of  oxidizing  sulphurous  acid  to  sulphuric  acid 
proceeds  ideally  according  to  the  reaction 

H2S03  +  I2  +  H20  =  H2SO4  +  2  HI. 

In  more  concentrated  solutions  of  sulphur  dioxide  the  reaction 
may  result  in  the  formation  of  dithionic  acid  instead  of  sulphuric 
acid  exclusively: 

2  H2S03  +  I2  =  H2S2O6  +  2  HI. 

Moreover,  a  reaction  may  set  in  between  the  sulphur  dioxide  and 
the  hydriodic  acid  produced  or  present  in  consequence  of  the 

*  See  p.  166. 


172  QUANTITATIVE  CHEMICAL  ANALYSIS 

action  of  the  sulphuric  acid  upon  the  potassium  iodide  of  the 
standard  iodine  solution.  In  this  secondary  reaction  iodine 
and  sulphur  are  set  free: 

S02  +  4  HI  =  2  I2  +  2  H20  +  S. 

When  the  conditions  of  reaction  are  so  adjusted  that  there  is 
no  excess  of  sulphur  dioxide  to  react  with  iodine  or  with  hydri- 
odic  acid  the  subsidiary  reactions  cannot  take  place;  and  this 
is  the  case  when  the  solution  of  sulphurous  acid  is  slowly  added 
to  the  iodine  solution  or  when  the  solution  of  a  sulphite  is 
slowly  added  to  the  iodine  solution  previously  acidulated.  In 
determining  sulphurous  acid  or  a  sulphite  the  usual  course  of 
the  iodine  titration  is  therefore  reversed;  the  solution  of  the 
substance  to  be  determined  is  made  up  to  a  known  volume, 
and  of  this  solution  a  measured  portion  of  the  standard  iodine, 
suitably  acidified,  is  added  to  the  point  of  decoloration.* 

Action  of  Iodine  in  Excess  and  Titration  of  the  Excess 

Determination  of  Soluble  Sulphides.  —  Iodine  reacts  with 
hydrogen  sulphide,  according  to  the  equation 

H2S  +  I2  =  2  HI  +  S, 

but  in  the  practical  determination  of  the  volatile  hydrogen 
sulphide  by  means  of  this  reaction,  direct  titration  by  gradual 
addition  of  standard  iodine  is  not  feasible.  The  process  may, 
however,  be  made  successful  by  introducing  the  solution  of 
hydrogen  sulphide  into  a  solution  of  standard  iodine,  so  di- 
luted that  the  separated  sulphur  will  not  coagulate  to  a  film, 
and  determining  by  standard  sodium  thiosulphate  the  excess  of 
the  latter.  Sulphides  and  sulphydrates  which  evolve 'hydrogen 
sulphide  if  brought  into  contact  with  dilute  hydrochloric  acid 
may  be  similarly  determined  by  introducing  them  into  the 
diluted  solution  of  standard  iodine  containing  also  hydro- 
chloric acid  —  or,  by  acting  upon  them  with  acid  and  conduct- 
ing the  evolved  hydrogen  sulphide  into  the  acidulated  and 

*  See,  also,  Ashley:  Methods  in  Analysis,  Gooch,  p.  366. 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS 


173 


dilute  solution  of  standard  iodine  —  the  iodine  in  excess  being 
titrated  by  standard  thiosulphate. 

Processes  Based  upon  Liberation  of  Iodine  in  Solution 

Determination  of  Bromine.  —  With  an  excess  of  potassium 
iodide,  bromine  reacts  according  to  the  equation 

Br2  +  2  KI  =  2  KBr  +  I2. 

When  bromine  in  water  solution  (best  faintly  acidulated  with 
hydrochloric  acid)  is  added  to  the  solution  of  potassium  iodide 
the  iodine  set  free  measures  the  bromine.  To  avoid  loss  by 
volatilization  in  the  process  the  bromine  solution  should  be  de- 
livered from  a  pipette,  the  opening  of  which  is  held  close  to  the 
surface  of  the  iodide  solution  contained  in  a  flask  or  bottle  pro- 
vided with  a  glass  stopper.  Upon  shaking  the  liquids  together 
in  the  stoppered  bottle  the  reaction  is  completed  immediately 
and  the  free  iodine,  which  may  be  titrated  by  standard  thio- 
sulphate or  with  standard  arsenite  after  the  addition  of  an 
excess  of  alkali  hydrogen  carbonate,  measures  the  bromine 
originally  free,  according  to  the  equation  above. 

When  the  bromine  is  in  gaseous  form,  as  when  it  is  dissolved 
in  air,  it  must  be  introduced  into  the  solution  of  potassium 
iodide  in  a  way  to  preclude  danger  of  mechanical 
loss  in  the  process  of  absorption  by  the  iodide 
solution. 

A  convenient  method  for  the  manipulation  of 
such  gaseous  substances  is  to  place  them,  con- 
tained in  a  glass  stoppered  pipette  or  Chancel 
flask,  between  a  carbon  dioxide  generator  and  a 
trapped  absorption  bottle,  of  the  form  shown  in 
Fig.  28,  charged  with  a  solution  of  potassium 
iodide.  The  gas  may  then  be  passed  bubble  by 
bubble  into  the  solution  of  potassium  iodide  which 
reacts  with  the  bromine  and  retains  most  of  the  iodine  set  free, 
while  any  traces  of  bromine  or  iodine  which  may  escape  from 
the  bottle  are  caught  in  the  trap  also  charged  with  potassium 
iodide. 


FIG.  28. 


174  QUANTITATIVE  CHEMICAL  ANALYSIS 

From  the  number  of  cubic  centimeters  (ri)  and  the  iodine  equivalent 
per  cm.3  (i)  of  the  thiosulphate  or  arsenite,  the  weight  of  bromine  (it?)  may 
be  calculated  from  the  proportion 

w  :  ni  =  Br  :  I  =  79.97  :  126.92, 
=  ni  X  79.97 
126.92 

Determination  of  Chlorine.  —  Chlorine  in  water  solution  or 
in  gaseous  form  may  be  determined  by  procedures  exactly  like 
those  described  for  the  determination  of  bromine,  the  chlorine 
reacting  with  an  excess  of  potassium  iodide  according  to  the 
equation 

C12  +  2  KI  =  2  KC1  +  I2, 

in  accordance  with  which  the  amount  of  chlorine  may  be  cal- 
culated from  the  iodine  liberated  and  determined  by  titration 
with  the  standard  thiosulphate  or  the  standard  arsenite.  « 

From  the  number  of  cubic  centimeters  (n)  and  the  iodine  equivalent 
per  cm.3  (i)  of  the  thiosulphate  or  arsenite,  the  weight  (w)  of  chlorine  is 
given  by  the  formula 

ni  X  35.45 


w  = 


126.92 


The  Reaction  between  an  Iodide  and  an  lodate  in  Presence 
of  Dilute  Acid.  —  From  iodides  the  iodine  may  be  liberated  by 
the  action  of  iodic  acid  or  of  an  iodate  and  a  suitable  acid,  e.g., 
dilute  sulphuric  acid;  from  iodates  (and  periodates)  the  iodine 
may  be  liberated  by  the  action  of  a  suitable  acid  and  an  iodide. 
In  either  process  iodic  acid  and  hydriodic  acid  enter  into  the 
reversible  reaction 

HI03  +  5  HI  <=»  3 12  +  3  H20. 

Under  suitable  conditions  this  reaction  may  be  utilized  for  the 
determination  of  iodine  in  iodides  or  in  iodates.  At  all  volumes 
ordinarily  used  in  analysis  the  direct  action  is  practically  com- 
plete when  the  hydriodic  acid  is  in  moderate  excess.  In  order 
that  the  reaction  given  above  may  be  realized  when  the  iodic 
acid  is  in  excess  attention  must  be  paid  to  limiting  the  volume 
of  the  water  solution  in  which  the  reaction  takes  place.  In 
presence  of  an  appreciable  amount  of  hydrochloric  acid,  or  a 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  175 

chloride,  a  secondary  reaction  takes  place  which  results  in  the 
formation  of  iodine  chloride;  and  this  action  may  be  com- 
plete in  presence  of  a  suitable  amount  of  hydrochloric  acid :  * 

2 12  +  KIO3  +  5  HC1  =  5 IC1  +  3  H20. 

Determination   of   lodates.  —  When   potassium   iodide   in 
excess  and  dilute  hydrochloric  acid,  or  dilute  sulphuric  acid, 
are  added  to  the  solution  of  an  iodate,  iodine  is  immediately 
set  free  according  to  the  equation 

MI03  +  5  KI  +  6  HC1  =  5  KC1  +  MCI  +  3  H20  +  3 12. 

The  iodine  set  free  may  be  titrated  at  once  by  standard  sodium 
thiosulphate,  or  it  may  be  absorbed  by  an  excess  of  standard 
arsenite,  this  excess  being  determined  by  titration  with  standard 
iodine.  In  these  operations 

MI03  *  3  I2. 

Experimental  Process:  Determination  of  Iodate  by  Thiosulphate  Titra- 
tion of  Iodine.  —  Weigh  out  exactly  about  0.15  grm.  of  potassium  iodate, 
dissolve  it  in  water  (about  100  cm.3),  add  in  solution  an  excess  of  potas- 
sium iodide  (1  grm.)  and  acidulate  the  mixture  with  dilute  hydrochloric 
acid  or  sulphuric  acid.  Titrate  at  once  the  free  iodine  by  means  of  N/10 
thiosulphate.  Find  the  weight  of  iodate  (IP)  from  the  number  of  cubic 
centimeters  (ri)  and  the  iodine  equivalent  per  cm.3  (i)  of  the  thiosulphate 
solution  by  the  proportion 

ni  :  w  =  6  X  126.92  :  1  grm.  mol.  MI03, 
_  ni  X  1  grm.  mol.  MIO3 
6  X  126.92 

Tabulate  the  details  of  the  determination. 

Experimental  Process:  Determination  of  Iodate  by  Arsenite  Absorption  of 
Iodine.  —  When  the  iodine  set  free  in  the  reaction  between  an]  iodate,  an 
iodide,  and  hydrochloric  acid  is  estimated  by  acting  upon  it  with  standard 
arsenite  in  excess  and  titrating  this  excess  by  standard  iodine,  precaution 
must  be  taken  to  prevent  any  escape  of  iodine  with  the  carbon  dioxide 
evolved  in  the  process  of  neutralization.  A  suitable  apparatus  for  the  ap- 
plication of  the  process  is  a  reaction-bottle  of  500  cm.3  or  1000  cm.3  capacity 
(Fig.  28,  p.  173)  provided  with  a  stopper  carrying  a  stop-cock  and  thistle 
fused  to  the  inlet  tube  and  a  Will  and  Varrentrapp  absorption  trap  sealed 
to  the  outlet  tube. 

*  See  p.  178. 


176  QUANTITATIVE  CHEMICAL  ANALYSIS 

Put  the  iodate  to  be  determined  (about  0.15  grm.  of  KIO3)  into  the 
bottle  with  more  than  the  equivalent  amount  of  potassium  iodide  (3  grm.) 
and  adjust  the  volume  of  the  liquid  to  about  150  cm.3.  Set  the  stopper 
in  place,  and  charge  the  trap  with  a  solution  (5  per  cent)  of  potassium 
iodide.  Introduce  sulphuric  acid  (5  cm.  of  the  [1  :  3]  mixture)  through  the 
thistle  tube  and  agitate  the  solution  thoroughly  by  a  rotary  motion  of  the 
bottle,  to  bring  about  the  mixture  of  the  acid,  the  iodate,  and  the  iodide. 
Add  an  amount  of  sodium  hydroxide  in  solution  sufficient  (as  previously 
determined)  to  neutralize  the  larger  part  of  the  free  acid,  shake  thoroughly, 
and  add  a  saturated  solution  of  potassium  hydrogen  carbonate,  KHCO3, 
to  an  amount  about  10  cm.3  more  than  sufficient  to  neutralize  all  the  sul- 
phuric acid.  Close  the  stop-cock  and  shake  the  solution  carefully  until  the 
last  traces  of  iodine  vapor  have  been  absorbed  by  the  liquid.  Run  the 
solution  in  the  trap  into  the  bottle  and  add  the  washings  of  trap,  tubes, 
and  stopper.  Now,  introduce  standard  N/10  arsenite  to  the  bleaching 
point  of  the  iodine,  add  starch  solution,  and  titrate  the  excess  of  arsenite 
with  N/10  iodine  (usually  only  a  few  drops)  to  coloration.  Calculate  the 
weight  of  the  iodate  (w)  from  the  number  of  cubic  centimeters  (n)  and  the 
iodine  equivalent  per  cm.3  (i)  =  (0.012692)  of  the  N/10  arsenite  solution, 
with  the  number  of  cubic  centimeters  of  the  standard  iodine  (n')  and  its 
iodine  value  per  cm.3  (i),  by  the  proportion 

ni  —  rii'  :  w  =  6  X  126.92  :  1  grm.  mol.  MIO3, 
=  (ni  -  n'i'}  X  1  grm.  mol.  MIO3 
6  X  126.92 

Tabulate  the  details  of  the  operation. 

Determination  of  Periodates.  —  Action  of  an  Iodide  in  Pres- 
ence of  Acid.  Periodates  may  be  analyzed  by  procedure  similar 
to  that  described  above  for  the  determination  of  iodates,  the  de- 
composition taking  place  according  to  the  equation 

MI04  +  7  KI  +  8  HC1  =  7  KC1  +  MCI  +  4 12. 

If  the  process  of  direct  titration  of  the  free  iodine  by  standard  thiosul- 
phate  is  followed,  the  calculation  may  be  made  from  the  number  of  cubic 
centimeters  (n)  of  the  thiosulphate  used  and  the  iodine  equivalent  per  cm.3 
(t)  of  that  reagent,  according  to 

_  ni  X  1  grm.  mol.  MIO4 
8  X  126.92 

If  the  titration  of  the  free  iodine  is  made  by  using  standard  arsenite  in 
excess  (n  cubic  centimers  of  iodine  equivalent  i),  and  estimating  the  excess 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          177 

by  standard  iodine  (ri  cubic  centimeters  of  iodine  value  i')  the  calculation 
may  be  made  by  the  formula 

=  (ni  -  nV)  X  1  grm.  mol.  MIO4 
8  X  126.92 

Action  of  an  Iodide  in  Alkaline  Solution.  The  decomposition 
of  a  periodate  may  also  be  effected  by  potassium  iodide  in  pres- 
ence of  an  alkali  hydrogen  carbonate,  according  to  the  equation 

MIO4  +  2  KI  +  2  MHC03  =  MI03  +  K2C03 

+  M2CO3  +  H20  +  I2. 

The  solution  of  the  periodate  is  made  barely  alkaline  to  phe- 
nolphthalein  with  the  use  of  an  alkali  hydroxide  or  hydrochloric 
acid  as  may  be  necessary,  and  is  then  treated  with  sodium  hy- 
drogen carbonate  (10  cm.3  of  the  cold  saturated  solution)  and 
potassium  iodide.  The  iodine  liberated  immediately  may  be  ti- 
trated by  standard  N/10  arsenite  which  is  not  affected  by  the 
iodate,  and  this  reaction  may  be  utilized  for  the  determination 
of  the  iodine  of  a  periodate  in  presence  of  an  iodate. 

If  n  cubic  centimeters  of  arsenite  having  an  iodine  equivalent  per  cm.3 
i  are  used,  the  weight  w  of  periodate  is  given  by  the  following  expression : 
=  n  X  0.012692  X  1  grm.  mol.  MIO4 
2  X  126.92 

Determination  of  lodates  and  Periodates  in  Association.  — 

The  weight  of  iodine  (w)  of  an  iodate  associated  with  the  perio- 
date may  be  determined  by  combining  the  result  of  the  titration 
of  the  iodine  set  free  (from  the  periodate)  in  the  alkaline  solution 
with  that  of  the  titration  of  the  iodine  (from  the  iodate  and 
periodate)  set  free  in  acid  solution. 

Denoting  by  n  the  number  of  cubic  centimeters  used  of  the  N/10 
arsenite  and  by  n'  the  number  of  cubic  centimeters  of  thiosulphate  con- 
taining the  iodine  equivalent  per  cm.3  if, 

=  (n'i'  -  4  [n  X  0.012692])  X  1  grm.  mol.  MIO3 
6  X  126.92 

Determination  of  Iodides.  —  Action  of  the  Iodate  in  Slightly 
Acid  Solution.  In  applying  the  reaction 

5  HI  +  HI03  ^  3  H20  +  3 12, 


178  QUANTITATIVE  CHEMICAL  ANALYSIS 

to  the  analysis  of  an  iodide,  it  is  obvious  that  if  the  iodine  set 
free  in  presence  of  an  excess  of  the  iodate  is  to  be  determined 
in  the  solution  and  made  the  measure  of  the  iodide  it  must  be 
determined  by  a  reagent  which  will  not  be  affected  also  by  the 
excess  of  iodate.  Sodium  thiosulphate  is  not  available  for  this 
purpose,  since  in  acid  solution  it  reacts  with  the  iodate  while  in 
alkaline  solution  its  reaction  with  iodine  is  irregular.  But  it 
is  possible  to  so  arrange  the  process  that  the  action  between  the 
iodate  and  the  iodide  takes  place  in  acid  solution  and  the  deter- 
mination of  the  liberated  iodine  is  made  in  alkaline  solution  by 
means  of  the  standard  arsenite,  which  is  unaffected  by  the  ex- 
cess of  the  iodate.*  The  reaction  between  an  iodide,  iodate  and 
acid  may  also  be  applied  in  such  fashion  that  the  difference  be- 
tween the  amount  of  iodate  taken  in  excess  and  the  amount  of 
it  which  remains  after  boiling  out  the  free  iodine  (see  p.  179) 
shall  be  made  the  measure  of  the  iodine  to  be  determined.! 

In  either  of  these  processes,  the  presence  of  any  considerable 
amount  of  chloride  or  bromide  interferes  with  the  accuracy  of 
the  analytical  process,  since  chlorides  and  bromides  are  ap- 
preciably attacked  by  iodic  acid  and  iodine  in  association,  with 
the  formation  of  iodine  chloride  and  iodine  bromide. 

Action  of  the  Iodate  in  Presence  of  Strong  Hydrochloric  Acid 

The  iodine  set  free  in  the  reaction  between  a  standard  solution 
of  iodate  and  an  iodide  (see  also  p.  180)  may  be  converted  to 
iodine  monochloride  by  further  action  of  the  iodate  in  presence 
of  sufficiently  concentrated  hydrochloric  acid,  and  the  amount 
of  iodine  acted  upon  may  be  found  from  the  amount  of  iodate 
required  to  complete  this  reaction,!  according  to  the  equation 

KI03  +  2  KI  +  6  HC1  =  3  KC1  +  3 IC1  +  3  H20. 

The  vanishing  of  the  iodine  color  from  chloroform  shaken  up 
with  this  solution  marks  the  end-point  of  the  reaction.  This 

*  Gooch  and  Walker:   Methods  in  Analysis,  Gooch,  p.  454. 
t  Dietz  and  Margosches:  Analytical  Chemistry,  Treadwell-Hall,  3rd  ed., 
Vol.  II,  p.  671. 

t  Andrews:  Jour.  Am.  Chem.,  25,    756  (1903). 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  179 

process  has  the  advantage  that  the  reagent  which  sets  free  the 
iodine  determines  it  also,  and  that  the  presence  of  chlorides  or 
bromides  does  not  interfere  with  the  accuracy  of  the  determi- 
nation. 

Experimental  Process:  Arsenite  Determination  of  Iodine  Set  Free  from 
Iodide  by  lodate.  —  Draw  from  a  burette  in  a  reaction  bottle  (Fig.  28), 
an  amount  of  an  approximately  N/10  solution  of  potassium  iodide  equiv- 
alent to  about  0.4  grm.  of  the  iodide.  Add  more  than  equivalent  amount 
(about  0.2  grm.)  of  potassium  iodate  in  solution  and  make  up  the  volume  of 
liquid  to  150  cm.3.  Insert  the  stopper  and  charge  the  trap  with  a  solution 
(5  per  cent)  of  potassium  iodide.  Introduce  sulphuric  acid  (5  cm.3  [1  :  3]) 
through  the  thistle  tube  and  mix  the  solution.  Then  add  slowly  a  satu- 
rated solution  of  potassium  hydrogen  carbonate,  in  amount  (10  cm.3) 
somewhat  in  excess  of  that  needed  to  neutralize  the  sulphuric  acid.  After 
the  neutralization,  which  prevents  interaction  when  the  iodide  in  the 
trap  and  the  iodate  in  the  bottle  come  into  contact,  bleach  the  free  iodine 
in  the  bottle  and  trap  by  N/10  arsenite,  add  starch,  and  titrate  with  N/10 
iodine  the  very  slight  excess  of  the  arsenite  used  in  the  bleaching.  From 
the  iodine  equivalent  per  cm.3  of  the  standard  arsenite  (i)  and  the  number 
of  cubic  centimeters  of  it  used  (n),  with  the  iodine  value  per  cm.3  of  the 
standard  iodine  (if),  and  the  number  of  cubic  centimeters  of  it  used  (n'}, 
calculate  the  weight  (w)  of  the  iodine  given  by  the  proportion 

ni  —  n'i'  :  w  =  3I2  :  SMI, 

_  5  (ni  —  n'i')  (1  grm.  mo.  MI) 
6  X  126.92 

Fix  the  standard  of  the  solution  of  iodide. 

Experimental  Process:  Determination  of  Iodide  from  Excess  of  lodate 
Used.  —  First,  make  a  solution  of  potassium  iodate  of  nearly  N/10  value  in 
respect  to  the  iodine  which  will  be  evolved  by  its  reaction  with  an  iodide  in 
presence  of  acid  (containing  in  a  liter  Igrm.  mol.  •*-  6  X  10  of  KIO3)  and  de- 
termine the  exact  iodine  equivalent  per  cubic  centimeter  by  treating  a 
measured  portion  with  potassium  iodide  in  excess,  adding  sulphuric  acid, 
and  titrating  by  N/10  thiosulphate  the  iodine  set  free.  From  the  iodine 
equivalent  per  cm.3  of  the  thiosulphate  (i)  and  the  number  of  cubic  centi- 
meters used  (n),  with  the  number  of  cubic  centimeters  (n'}  of  the  iodate, 
the  iodine  equivalent  per  cm.3  of  the  iodate  (i'}  solution  will  be  given  by  the 
expression 


Next,  treat  the  solution  of  the  iodide  to  be  determined  with  a  measured 
volume  of  the  iodate  solution   (now  of  known  iodine  value  per  cm.3), 


180  QUANTITATIVE  CHEMICAL  ANALYSIS 

acidify  with  hydrochloric  acid,  introduce  a  piece  of  calcite,  and  boil 
until  the  free  iodine  is  expelled.  In  this  process  the  iodide  is  completely 
broken  up  in  accordance  with  the  equation 

KIO3  +  5  MI  +  6  HC1  =  5  KC1  +  MCI  +  3  H2O  +  3  12, 

the  amount  of  iodide  to  be  determined  being  equivalent  to  five-sixths  of 
the  iodate  which  disappears.  Find  the  amount  of  the  iodate  which  re- 
mains by  treating  the  cooled  solution  with  an  excess  of  potassium  iodide 
and  titrating  with  the  standard  thiosulphate  the  iodine  liberated,  as  in 
the  determination  of  the  iodine  of  the  iodate. 

If  the  iodine  equivalent  per  cm.3  of  the  thiosulphate  is  denoted  by  i,  the 
number  of  cubic  centimeters  used  in  titrating  the  iodine  liberated  by  the  ex- 
cess of  iodate  by  n,  the  iodine  equivalent  per  cm.3  of  the  iodate  by  i'  ,  and 
the  number  of  cubic  centimeters  of  iodate  used  by  n',  the  amount  of  iodine 
liberated  by  the  iodate  which  disappeared  in  reaction  with  the  iodide 
to  be  determined  will  be  given  by  the  expression  n'i'  —  ni,  and  five-sixths  of 
this  amount  will  be  the  iodine  equivalent  of  that  iodide.  The  weight  of 
iodide  determined  is  given  by  the  proportion 

5/6  (n'i'  -ni)  :w  =  l:Ml  =  126.92  :  1  grm.  mol.  MI, 
=  5  (riir  -ni)  (1  grm.  mol.  MI) 
6  X  126.92 

Tabulate  the  details  of  the  operation. 

Experimental  Process:  Determination  of  Iodide  by  Direct  Action  of 
Standard  Iodate.  —  Draw  from  a  burette  20  cm.3  of  the  approximately 
N/10  solution  of  potassium  iodide  into  a  250-cm.3  glass  bottle  provided 
with  a  glass  stopper.  Add  5  cm.3  of  chloroform  and  30  cm.3  of  concentrated 
hydrochloric  acid.  Into  the  mixture  measure  from  a  burette  successive 
portions  of  the  standard  iodate  solution,  inserting  the  stopper  and  shaking 
the  contents  of  the  bottle  vigorously  between  the  additions  of  the  iodate, 
until  the  violet  color  of  iodine  in  the  chloroform  just  disappears.  Cal- 
culate the  standard  of  the  iodide  solution  from  the  relation 


Other  Substances  (Oxidizers)  Reactive  with  Iodide  or  Hydri- 
odic  Acid.  —  Like  the  iodates  and  periodates,  certain  other 
substances  enter  into  complete  reactions  with  potassium  iodide 
or  with  hydriodic  acid  (produced  in  the  interaction  of  an  excess 
of  potassium  iodide  and  a  suitable  acid),  so  that  exact  deter- 
minations of  these  substances  may  be  based  upon  the  im- 
mediate titration  of  the  iodine  thus  set  free.  Examples  of  such 

*  See  p.  178. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          181 

substances  are  hydrogen  dioxide,  the  soluble  hypochlorites, 
bromates,  chromates,  permanganates,  the  higher  oxides  of  lead 
and  cerium,  gold  tri-chloride,  ferric  chloride,  and  cupric  salts. 

Determination  of  Hydrogen  Dioxide.  —  In  dilute   solution 
and  in  presence  of  free  acid,  hydrogen  dioxide  liberates  iodine 
from  potassium  iodide  according  to  the  equation 
2HI-{-H202  =  2H2O  +  I2. 

If  a  suitable  amount  of  hydrogen  dioxide  (10  cm.3  of  a  0.6 
per  cent  solution)  is  added  gradually  to  dilute  (1 : 20)  sulphuric 
(200  cm.3)  containing  an  excess  (2  grm.)  of  potassium  iodide, 
the  iodine  set  free  after  a  short  interval  (5  minutes)  may  be 
titrated  with  N/10  thiosulphate  and  taken  as  the  measure  of 
the  hydrogen  dioxide. 

Determination  of  Hypochlorites.*  --  The  reaction  between 
hypochlorous  acid  and  hydriodic  acid  is  immediate,  so  that  in 
the  determination  of  a  hypochlorite  it  is  only  necessary  to  add 
it  to  a  solution  of  potassium  iodide,  acidify  the  hydrochloric 
acid,  and  titrate  the  free  iodine  with  sodium  thiosulphate. 
HOC1  +  HC1  +  2  KI  =  2  KC1  +  H2O  +  I2. 

Determination  of  Bromates.  —  Bromic  acid  and  hydriodic 
acid,  formed  in  the  interaction  of  potassium  iodide,  sulphuric 
acid,  and  potassium  bromate,  react  together  according  to  the 
equation 

6  HI  +  HBrO3  =  HBr  +  3  I2. 

Time  of  action,  proportion  of  iodide  to  bromate,  excess  of  acid, 
and  the  concentrations  are  all  within  limits  determining  factors, 

*  In  presence  of  a  chlorate  the  determination  of  the  hypochlorite  is 
best  made  otherwise,  since  the  former  may  take  part  to  some  extent  in 
the  evolution  of  iodine  from  the  iodide.  By  the  action  of  standard  arsenite 
the  hypochlorite  may  be  reduced  while  the  chlorate  is  not  affected: 

2  NaOCl  +  AsaOa  =  AsaCX  +  2  NaCl. 

In  the  technical  analysis  of  "  chloride  of  lime,"  which  may  also  contain 
chlorate,  it  is  customary  to  titrate  the  emulsion  of  this  substance  in  water 
with  standard  arsenite  to  the  point  at  which  a  single  drop  of  the  solution, 
removed  on  a  rod,  fails  to  set  free  iodine  and  produce  a  blue  coloration 
of  test-paper  made  by  impregnating  it  with  starch  and  potassium  iodide 
and  drying.  In  this  process  the  reaction  involving  iodine  is  only  employed 
to  indicate  the  end-point. 


182  QUANTITATIVE  CHEMICAL  ANALYSIS 

but  if  free  acid  and  a  considerable  excess  of  potassium  iodide 
are  present  the  reaction  proceeds  to  completion  within  a  reason- 
able time.  When  an  amount  of  potassium  bromate  not  exceed- 
ing 0.14  grm.  in  100  cm.3  of  solution  is  put  in  the  reaction 
bottle  (Fig.  28),  with  2.5  cm.3  of  sulphuric  acid  (1  :  1),  or  an 
equivalent  amount  of  hydrochloric  acid,  and  about  four  times 
the  amount  of  iodide  theoretically  required  (3  grm.)  the  re- 
action is  complete  in  half  an  hour  and  the  determination  of  the 
liberated  iodine  serves  to  measure  the  bromate  quite  exactly.* 
Determination  of  Chromates.  —  Chromic  acid  is  immedi- 
ately reduced  by  potassium  iodide  in  presence  of  hydrochloric 
acid,  to  the  green  chromic  chloride,  with  evolution  of  an  amount 
of  iodine  which  is  equivalent  to  the  oxygen  lost  by  the  chromic 
acid.  Chromates  and  dichromates  behave  similarly: 

2  K2Cr04  +  6  KI  +  16  HC1  =  10  KC1  +  2  CrCl3  +  8  H20  +*3  I2, 
K2Cr207  +  6  KI  +  14  HC1  =    8  KC1  +  2  CrCl3  +  7  H20  +  3 12. 

The  iodine  set  free  is  a  measure  of  the  chromate  or  dichromate 
entering  into  the  reaction.  After  suitably  diluting  the  solution, 
so  that  the  green  color  may  not  mask  the  disappearance  of  the 
starch  blue  in  the  end  reaction,  the  greater  part  of  the  iodine 
is  bleached  by  standard  thiosulphate,  starch  is  added,  and  the 
titration  by  the  thiosulphate  is  continued  until  the  color  of  the 
solution  changes  from  blue  to  green. 

Permanganates.  —  Upon  contact  with  a  solution  of  potassium 
iodide  containing  hydrochloric  acid,  permanganates  are  re- 
duced with  the  formation  of  manganous  chloride  and  the 
liberation  of  iodine: 

2  KMn04  + 10  KI  +  16  HC1  =  12  KC1  +  2  MnCl2  +  8  H20  +  5I2. 

The  iodine  liberated,  which  is  the  measure  of  the  permanganate 
originally  present,  may  be  titrated  by  the  process  previously 
described  for  the  standardization  of  the  thiosulphate  solution 
by  means  of  potassium  permanganate.  (See  p.  167.) 

Determination  of  Lead  Dioxide.  —  In  presence  of  5  per  cent 
acetic  acid,  sodium  acetate,  and  potassium  iodide,  lead  dioxide, 

*  Gooch  and  Blake:   Methods  in  Analysis,  Gooch,  p.  471. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          183 

finely  ground  (if  dry)  or  in  hydrated  form,  is  decomposed  ac- 
cording to  the  equation 

PbO2  +  4  HI  =  PbI2  +  2  H20  +  I2. 

If  a  large  excess  of  the  acetate  is  present  with  not  too  much 
potassium  iodide  the  lead  iodide  will  not  precipitate,  and  upon 
further  dilution  the  dissolved  iodine  may  be  titrated  with  stand- 
ard thiosulphate. 

Suitable  proportions  are  about  0.5  grm.  of  lead  dioxide, 
1.2  grm.  of  potassium  iodide,  10  grm.  of  sodium  acetate,  and 
5  cm.3  of  5  per  cent  acetic  acid.  If  lead  iodide  is  precipitated 
it  should  be  dissolved  by  more  sodium  acetate  with  a  little 
water.  After  dilution  to  a  volume  of  about  25  cm.3,  the  still 
clear  solution  of  iodine  is  titrated  with  the  thiosulphate. 

Determination  of  Cerium  Dioxide.  —  Anhydrous  cerium 
dioxide,  prepared  by  the  ignition  of  the  oxalate  or  the  hydrox- 
ide, is  very  slowly  acted  on  by  hydrochloric  acid,  especially 
when  free  from  the  oxides  of  lanthanum  and  didymium.  By 
digesting  cerium  dioxide  with  potassium  iodide  and  hydro- 
chloric acid  the  decomposition  of  the  dioxide  is  brought  about 
and  iodine  is  set  free  in  accordance  with  the  reaction 

2Ce02  +  8HC1  +  2KI  =  2CeCl3  +  2KC1  +  4H20  '+  I2. 

In  this  procedure,  a  weighed  portion  of  the  cerium  dioxide  is 
placed  in  a  small  glass-stoppered  bottle  (of  about  100  cm.3 
capacity),  with  1  grm.  of  potassium  iodide  and  a  few  drops  of 
water.  Carbon  dioxide  is  passed  into  the  bottle  to  expel  the 
air,  concentrated  hydrochloric  acid  (10  cm.3)  is  added,  the 
stopper  is  inserted,  and  the  bottle  is  heated  gently  upon  a 
steam  radiator  for  about  an  hour,  until  the  dioxide  is  com- 
pletely dissolved.  After  cooling,  the  solution  in  the  bottle  is 
carefully  washed  into  a  considerable  volume  of  water  and 
titrated  with  N/10  thiosulphate.  Correction  should  be  made 
for  the  amount  of  iodine  set  free  from  potassium  iodide  by 
hydrochloric  acid,  without  the  addition  of  cerium  dioxide,  un- 
der similar  conditions  of  treatment,  but  this  correction  should 
not  exceed  a  few  hundredths  of  a  cubic  centimeter  if  the  process 
is  carefully  conducted. 

Determination  of  Gold  Trichloride.  —  When  potassium 
iodide  reacts  at  suitable  concentrations  with  minute  amounts 


184  QUANTITATIVE  CHEMICAL  ANALYSIS 

of  gold  trichloride  in  solution,  the  reaction  takes  place  regularly 
along  the  line  of  the  following  equation : 

AuCl3+  2  KI  =  AuCl  +  2  KC1  +  I. 

The  iodine  liberated  may  be  determined  by  N/100  or  N/1000 
sodium  thiosulphate  and  the  amount  of  the  gold  may  be  cal- 
culated from  the  proportion  of  the  equation  given.*  In  a 
similar  process  in  which  the  reduction  of  gold  chloride  is  effected 
by  potassium  iodide  and  the  free  iodine  titrated  with  sulphur- 
ous acid,  larger  amounts  of  gold  may  be  determined  (Lenher). 
Determination  of  Ferric  Chloride.  —  The  reaction  between 
ferric  chloride  and  hydriodic  acid  is  reversible: 

2  FeCl3  +  2  HI  <=±  2  FeCl2  +  2  HC1  +  I2. 

In  order  that  the  direct  action  between  ferric  chloride,  potassium 
iodide,  and  hydrochloric  acid  may^  proceed  to  completion  it  is 
necessary  to  have  a  considerable  excess  of  the  iodide  present 
and  to  remove  from  time  to  time  the  iodine  set  free  while  pro- 
tecting the  acid  solution  of  the  iodide  from  action  of  atmos- 
pheric oxygen  during  the  protracted  titration. 

The  solution  of  ferric  chloride,  or  ferric  sulphate,  containing 
a  slight  excess  of  free  hydrochloric  acid  is  placed  in  a  trapped 
reaction-bottle  like  that  shown  in  Fig.  28,  p.  173.  The  stopper 
carrying  the  inlet  tube  with  thistle  and  stop-cock  and  the  outlet 
tube  with  its  trap  charged  with  a  solution  of  potassium  iodide 
is  put  in  place;  the  air  in  the  bottle  is  replaced  by  carbon 
dioxide  introduced  through  the  thistle  tube;  a  solution  of 
potassium  iodide  (5  grm.)  is  introduced  through  the  thistle  tube 
and  mixed  with  the  ferric  chloride  by  a  rotary  movement  of 
the  bottle;  and  the  mixture  is  allowed  to  stand  for  twenty 
minutes.  The  liberated  iodine  is  nearly  bleached  by  titration 
with  sodium  thiosulphate  introduced  through  the  thistle  tube, 
starch  is  added,  and  the  titration  is  continued  until  the  starch- 
iodide  blue  disappears.  If  upon  further  standing  the  blue 
color  returns,  more  iodine  must  be  added,  and  the  process  of 
standing  and  titrating  repeated. 

Determination  of  Copper  in  Cupric  Salts.  —  When  potas- 
sium iodide  is  added  to  the  solution  of  a  cupric  salt,  cuprous 
*  Gooch  and  Morley:  Methods  in  Analysis,  Gooch,  p.  146. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  185 

iodide  is  precipitated,  while  iodine  is  liberated  in  an  amount 
which  is  equivalent  to  the  cuprous  iodide: 

2  CuSO4  +  4  KI  =  2  K2SO4  +  Cu2I2  +  I2. 

This  reaction  has  been  made  the  basis  of  an  iodometric  method 
for  the  determination  of  copper.*  To  obtain  the  best  results, 
attention  must  be  paid  to  the  concentrations  of  the  reagents 
and  of  free  acid. 

The  procedure  best  suited  to  the  determination  of  an  amount 
of  copper  not  exceeding  0.3  grm.  seems  to  be  covered  by  the 
following  directions.!  The  solution  of  the  cupric  salt,  con- 
taining no  more  than  3  cm.3  of  concentrated  sulphuric  acid, 
hydrochloric  acid,  nitric  acid  (free  from  nitrogen  oxides)  or 
25  cm.3  of  50  per  cent  acetic  acid,  is  made  up  to  a  volume  of 
100  cm.3  and  treated  with  5  grm.  of  iodate-free  potassium 
iodide.  The  iodine  liberated  is  titrated  by  sodium  thiosulphate 
with  the  use  of  starch  indicator  at  the  end  of  the  titration. 
When  as  much  as  25  cm.3  of  the  N/10  sodium  thiosulphate  has 
been  used  in  the  titration  without  bleaching  all  the  iodine 
2  grm.  or  3  grm.  more  of  potassium  iodide  should  be  added  be- 
fore continuing  the  titration. 

Titration  of  Iodine  after  Extraction  from  Aqueous  Solution 

When  iodine  is  liberated  in  association  with  substances  which 
interfere  with  the  direct  titration  by  means  of  sodium  thiosul- 
phate or  by  an  alkali  arsenite,  it  is  sometimes  possible  to  esti- 
mate the  iodine  by  titration  after  extracting  it  from  the 
aqueous  solution  in  a  non-miscible  solvent,  such  as  carbon  di- 
sulphide  or  chloroform.  In  applying  the  process  of  extraction 
the  general  principle  which  governs  the  division  of  a  solute 
between  two  non-miscible  solvents  should  be  borne  in  mind.t 
The  experimentally  determined  ratio  of  the  concentration  per 
cm.3  of  the  iodine  in  water  to  that  in  the  carbon  disulphide  is 

*  Originally  proposed  by  DeHaen,  discussed  by  many  investigators, 

modified  and  advocated  by  Low.     For  review  and  references,  see  Methods 

.  in  Analysis,  Gooch,  p.  118:    for  the  application  to  the  technical  analysis 

of  copper  ores,  see  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed.,  pp. 

682-684. 

t  Gooch  and  Heath:  Methods  in  Analysis,  Gooch,  p.  121. 

t  See  p.  56. 


186  QUANTITATIVE  CHEMICAL  ANALYSIS 

1  :  400.  For  the  distribution  of  iodine  between  water  and 
carbon  disulphide,  the  general  expression,  in  which  XQ  is  the 
amount  of  iodine  in  the  water  and  xn  the  amount  of  iodine  in 
the  carbon  disulphide  after  n  shakings  of  the  water  solution  of 
volume  V  with  n  portions  of  carbon  disulphide  of  volume  v, 
takes  the  form 


° 


The  amounts  of  iodine  remaining  from  0.0500  grm.  after  shak- 
ing its  solution  in  100  cm.3  of  water  with  5  cm.3  portions  of 
carbon  disulphide  will  be  given  by  the  expression 


After  one  shaking  the  amount  of  iodine  remaining  in  thet  water 
separated  by  means  of  a  stoppered  funnel  will  be  0.0024  grm.; 
after  two  shakings,  0.0001  grm. 

Decomposition  of  Iodides  by  Nitrous  Acid.  —  The  deter- 
mination of  iodine  liberated  from  iodides  by  the  action  of 
nitrous  acid, 

2  HI  +  HN02  =  2  H20  +  NO  +  I2, 

and  titrated  with  sodium  thiosulphate  after  absorption  by 
shaking  the  aqueous  mixture  with  carbon  disulphide  or  with 
chloroform  rests  upon  this  principle.  This  process  is  subject  to 
irregularities  which  preclude  the  use  of  a  solution  of  the  thio- 
sulphate standardized  in  the  ordinary  way;  but  if  use  be  made 
of  an  empirical  standard  determined  by  matching  the  thio- 
sulphate against  the  iodine  set  free  by  nitrous  acid  from  a 
known  amount  of  iodide  closely  like  that  of  the  iodide  to  be 
determined,  the  process  may  be  made  serviceable  in  the  de- 
termination of  small  amounts  of  iodides  in  association  with 
relatively  large  amounts  of  chlorides  and  bromides. 

Processes   Depending   upon   the    Determination   of   Iodine 
Volatilized  to  the  Distillate 

The  non-reversible  reactions  in  which  hydriodic  acid  or  an 
iodide  is  completely  broken  up  with  evolution  of  iodine  by 
oxidizers  generally  involve  a  further  reaction  in  which  a  portion 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS 


187 


of  the  iodine  is  oxidized  to  iodic  acid.  To  remove  the  iodine 
completely  from  an  iodide  and  leave  it  free  so  that  it  may  be 
volatilized  to  the  distillate  it  is  necessary  to  employ  a  reaction 
in  which  a  product  of  reduction  is  sensitive  enough  to  oxidation 
to  prevent  the  formation  of  iodic  acid,  and  such  reactions  prove 
to  be  reversible  in  respect  to  this  liberation  of  iodine  from  an 
iodide.  When  the  iodine  evolved  is  made  the  measure  of  the 
oxidizer,  it  is  essential  to  exclude  oxygen  from  the  boiling  solu- 
tion containing  hydriodic  acid  in  order  that  iodine  may  not 
be  set  free  apart  from  the  main  reaction,  and  this  is  accom- 
plished by  performing  the  operation  in  an  atmosphere  of 
carbon  dioxide  made  as  free  as  possible  from  air  by  acting  in 
the  Kipp  generator  upon  previously  boiled  marble  with  boiled 
hydrochloric  acid  charged  with  cuprous  chloride. 

Determination  of  Iodine  in  Iodides.  —  Arsenic  acid  is 
capable  of  expelling  iodine  from  hydriodic  acid  without  forming 
iodic  acid.  The  reversible  reaction  between  arsenic  acid  in  ex- 
cess and  hydriodic  acid, 

H3As04  +  2  HI  <±  H3AsO3  +  H2O  +  I2, 

may  be  made  complete  in  the  forward  direction  by  volatilizing 
the  iodine,  and  if  the  iodine 
volatilized  be  absorbed  in 
the  distillate  it  maybe  easily 
determined  by  titration  and 
made  the  measure  of  the 
iodide  acted  upon. 

A  suitable  form  of  dis- 
tillation apparatus,*  con- 
structed with  sealed  or 
ground  joints  of  glass  where- 
ever  contact  with  iodine  is 
a  possibility,  as  shown  in 
Fig.  29,  may  be  easily  made  FIG.  29. 

by  sealing  together  a  sepa- 
rating funnel  A,  a  Voit   flask  B,  a  Drechsel  wash-bottle  C, 
and  a  bulbed  trap  g.     Upon  the  side  of  the  distillation  flask  B 
is  pasted  or  etched  a  graduated  scale  by  means  of  which  the 
*  Gooch  and  Norton:  Methods  in  Analysis,  Gooch,  p.  418. 


188  QUANTITATIVE  CHEMICAL  ANALYSIS 

volume  of  the  liquid  in  the  flask  may  be  known.  The  bottle 
used  as  the  receiver  and  the  trap  are  charged  with  a  solution 
of  potassium  iodide  (3  grm.  of  iodide  in  200  cm.3  of  water). 
The  iodide,  in  solution,  is  placed  in  the  distillation  flask  and 
the  connection  of  the  latter  with  the  receiver  is  adjusted.  A 
mixture  of  20  cm.3  of  sulphuric  acid  (1:1)  and  2  grm.  of  dihy- 
drogen  potassium  arsenate  is  introduced  through  the  stoppered 
funnel  followed  by  enough  water  to  make  up  the  total  volume 
of  liquid  to  about  50  cm.3.  Carbon  dioxide,  produced  in  a  Kipp 
generator  by  the  action  of  hydrochloric  acid  upon  marble,  is 
passed  through  the  apparatus,  to  facilitate  the  transfer  of  the 
iodine  from  the  distillation  flask  to  the  receiver.  The  liquid  is 
boiled  until  the  volume  has  decreased  to  35  cm.3,  when  the 
operation  should  be  stopped  and  the  iodine  in  the  receiver  and 
trap  titrated  by  sodium  thiosulphate  or,  after  the  addition  of 
an  excess  of  hydrogen  potassium  carbonate,  with  the  standard 
arsenite. 

Approximative  Determinations  of  Arsenic,  Antimony,  Vana- 
dium, Molybdenum,  Cerium,  and  Selenium.  —  Incomplete 
and  reversible  reactions  between  certain  oxidizers  and  potassium 
iodide  in  excess,  in  the  presence  of  an  acid,  may  also  be  made 
complete  by  the  removal  of  the  iodine  which  appears  as  a 
volatile  product  of  the  reaction.  Determinations  based  upon 
such  procedure  may  be  made  to  depend  upon  finding  the  amount 
of  iodine  distilled  and  collected  in  the  distillate;  but,  in  this 
case,  it  is  necessary  to  carry  on  the  process  of  distillation 
in  an  atmosphere  of  carbon  dioxide,  so  that  the  hydriodic  acid 
may  not  be  acted  upon  by  atmospheric  oxygen  with  the  con- 
sequent liberation  of  iodine  outside  the  main  reaction.  Fur- 
thermore, solutions  containing  hydriodic  acid  tend  to  withhold 
iodine  from  volatilization,  and  that  in  proportion  to  the  con- 
centration of  this  acid.  If  the  concentration  is  considerable  it 
is  impracticable  to  remove  to  the  distillate  the  last  traces  of 
iodine  set  free  by  an  oxidizer  from  an  iodide  used  in  excess. 
In  such  cases,  a  determination  of  the  oxidizer  must  generally 
be  made  to  depend  upon  determining  the  free  iodine  in  the 
residue  as  well  as  in  the  distillate.  In  this  manner  determina- 
tions may  be  made  of  arsenic,  antimony,  vanadium,  molybde- 
num, and  cerium,  in  compounds  showing  the  highest  valence 
of  these  elements,  and  of  selenium  in  selenious  and  selenic 
acids.  These  processes,  which  involve  the  distillation  and  de- 
termination of  the  iodine  liberated  in  the  interaction  of  the 
compounds  with  potassium  iodide  in  presence  of  hydrochloric 
acid,  are  based  upon  the  following  reactions: 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  189 

As2O5  +  4  HI  =  As2O3  +  2  H20  +  2  I2, 
Sb2O5  +  4  HI  =  Sb2O3  +  2  H2O  +  2  I2, 

V2O5  +  2  HI  =  V2O4  +  H2O  +  I2, 

SeO2  +  4  HI  =  Se  +  2  H2O  +  2 12, 

SeO3  +  6  HI  =  Se  +  3  H2O  +  3  I2, 

V2O5  -f  4  HI  =  V2O3  +  2  H2O  +  2 12, 

and,  at  extreme  concentration, 

2  Mo03  +  4  HI  =  2  Mo02I  =  2  H20  +  I2, 
2  CeO2  +  2  HI  =  Ce2O3  +  H2O  +  I2. 

The  Volatilization  of  Iodine  in  Acid  Solution  and  the 

Reoxidation   of   the   Residue   by   Iodine   in 

Alkaline    Solution 

Errors  due  to  the  action  of  atmospheric  oxygen  may  often 
be  eliminated  if  the  determination  is  based  upon  finding  the 
amount  of  the  reduced  product  in  the  residue. 

Determination  of  Iodides  by  the  Action  of  Arsenic  Acid  and 
Reoxidation  by  Iodine.  —  The  iodine  of  soluble  iodides  may 
be  accurately  determined  by  a  process  which  depends  upon 
breaking  up  the  iodide  by  the  action  of  an  arsenate  and  sul- 
phuric acid,  the  volatilization  of  the  iodine,  and  the  iodometric  * 
determination  of  the  arsenious  acid  retained  in  the  residue. 
The  reduction  proceeds  in  the  acid  solution  according  to  the 
expression 

H3  As04  +  2  HI  <=>  H3As03  +  H20  +  I2, 

and  is  completed  by  boiling  out  the  iodine.  The  arsenious  acid 
of  the  residue  is  determined,  after  neutralization,  by  titration, 
with  iodine  which  should  be  exactly  equivalent  in  amount  to 
iodine  evolved  from  the  iodide.  Under  conditions  to  be  de- 
scribed, the  presence  of  as  much  as  0.5  grm.  of  sodium  chloride 
does  not  increase  the  amount  of  arsenious  oxide  in  the  residue, 
but  does  induce  a  slight  loss  of  arsenious  oxide  (due  to  volatili- 
zation as  arsenious  chloride)  which  is  proportionate  to  the 
amount  of  the  last  and  the  amount  of  the  chloride.  The  effect 
of  a  bromide  is  to  produce  a  trifling  reduction  of  the  arsenic 

*  Gooch  and  Browning:  Methods  in  Analysis,  Gooch,  p.  457. 


190  QUANTITATIVE  CHEMICAL  ANALYSIS 

acid  without  volatility.  Due  correction  of  the  amount  of 
iodine  indicated  by  the  determination  of  the  arsenious  oxide  in 
the  residue  may  be  made  by  adding  to  the  indicated  amount 
0.008  of  the  product  of  the  weight  of  iodine  by  the  weight  of 
chlorine  in  the  chloride  and  subtracting  0.0024  of  the  weight 
of  bromine  in  the  bromide. 

Experimental  Process:  Determinations  of  Iodine  in  an  Iodide.  —  Dis- 
solve in  a  little  water  contained  in  an  Erlenmeyer  beaker  of  300  cm.3 
capacity,  the  substance  (which  should  not  contain  of  chloride  more  than 
the  amount  corresponding  to  0.5  grm.  of  sodium  chloride,  nor  of  bromide 
more  than  corresponds  to  0.5  grm.  of  potassium  bromide,  nor  of  iodide 
much  more  than  the  equivalent  of  0.5  grm.  of  potassium  iodide),  and  to 
the  solution  add  2  grm.  of  dihydrogen  potassium  arsenate 
dissolved  in  water,  20  cm.3  of  a  mixture  of  sulphuric  acid 
and  water  in  equal  volumes  (the  (1:1)  acid),  and  enough 
water  to  increase  the  total  volume  to  100  cm.3  or  a  little 
more.  Introduce  a  hollow  tetrahedron  of  thin  platinum 
foil  to  prevent  bumping  in  the  boiling,  hang  in  the  neck  of 
the  flask  a  trap  made  of  a  straight  two-bulb  drying  tube  cut 
short,  as  shown  in  Fig.  30,  and  boil  the  liquid  until  the 
level  reaches  a  mark  put  upon  the  flask  to  indicate  a  volume 
of  35  cm.3.  When  chlorides  and  bromides  are  present  great 
care  should  be  taken  not  to  press  the  concentration  beyond 
this  point  on  account  of  the  double  danger  of  losing  arsenious  chloride  and  of 
setting  up  reduction  of  the  arsenate  by  the  bromide.  On  the  other  hand, 
though  35  cm.3  is  an  ideal  volume,  failure  to  concentrate  below  40  cm.3 
introduces  no  appreciable  error.  Cool  the  liquid  remaining,  nearly  neutral- 
ize with  alkali  hydroxide  (ammonia  is  not  equally  good).  Complete  the 
neutralization  with  alkali  hydrogen  carbonate,  add  an  excess  of  20  cm.3  of 
the  saturated  solution  of  the  latter,  and  titrate  the  arsenite  in  solution 
with  N/10  iodine,  using  starch  as  the  indicator. 

From  the  number  of  cubic  centimeters  (n)  of  the  iodine  solution  used 
in  the  titration  and  the  iodine  value  per  cm.3  (i),  find  the  weight  of  iodine 
liberated  from  the  iodide  by  means  of  the  formula 

w  =  ni. 

Errors  due  to  the  presence  of  chlorides  and  bromides  may  be  eliminated 
by  adding  (algebraically)  to  the  iodine  indicated  (which  is  equal  to  the 
amount  of  iodine  used  in  the  process  of  titration),  an  amount 
i'  =  (0.008  X  wt.  Cl  X  wt.  I)  -  0.0024  X  wt.  Br. 

Determination  of  Arsenic  in  Arsenic  Acid.  —  The  deter- 
mination of  iodine  in  haloid  salts  based  upon  the  action  of 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  191 

arsenic  acid  in  presence  of  sulphuric  acid,  according  to  the 
equation 

H3As04  +  2  HI  =  H3As03  +  H2O  +  I2, 

the  iodine  being  completely  volatilized,  but  leaving  behind  in 
the  amount  of  arsenious  acid  produced  by  the  action  the  record 
of  the  iodine  of  the  iodide,  has  been  described  in  the  preceding 
paragraph.  In  the  determination  of  arsenic  in  arsenic  acid 
and  its  salts  the  reaction  is  utilized  conversely,  and  potassium 
iodide  in  excess,  in  presence  of  sulphuric  acid,  is  employed  to 
bring  about  the  reduction  of  arsenic  acid  to  arsenious  acid, 
which,  after  neutralization,  may  be  determined  by  titration 
with  standard  iodine.*  The  conditions  of  the  methods  are 
different,  in  that  in  the  former  the  hydriodic  acid  is  entirely 
broken  up  by  the  action  of  the  arsenic  acid,  with  the  complete 
volatilization  of  the  iodine,  while  in  the  latter  some  hydriodic 
acid  remains,  retaining  iodine  and  evolving  more  under  the 
action  of  the  air,  until  a  very  low  degree  of  concentration  is 
reached.  If  the  concentration  is  pushed  too  far  arsenious 
iodide  separates  from  the  sulphuric  acid  solution  and  then 
volatilizes.  The  small  amount  of  iodine  which  remains  dis- 
solved after  the  arsenic  acid  has  been  fully  reduced  may,  how- 
ever, be  converted  to  hydriodic  acid  by  the  careful  addition  of 
very  dilute  sulphurous  acid  (approximately  N/ 100),  and  if  cold 
water  is  added  as  soon  as  the  sulphurous  acid  has  bleached 
the  iodine,  and 'the  solution  immediately  neutralized,  the  return 
of  the  iodine  color  and  the  reversion  of  the  arsenious  acid  to 
arsenic  acid  are  precluded. 

Experimental  Process:  Determination  of  Arsenic  in  Arsenic  Acid.  — 
Heat  in  a  trapped  Erlenmeyer  flask  (see  Fig.  30)  the  solution  containing 
in  a  volume  of  50  cm.3  to  75  cm.3  an  amount  of  arsenate  not  more  than  the 
equivalent  of  0.4  grm.  of  arsenic  pentoxide,  with  potassium  iodide  to  an 
amount  about  0.5  grm.  in  excess  of  that  theoretically  required,  and  10  cm.3 
of  sulphuric  acid  (1  :  1).  After  introducing  folded  platinum  foil  to  prevent 
bumping,  boil  the  liquid,  taking  care  that  the  volume  never  becomes  less 
than  40  cm.3,  until  iodine  vapor  is  no  longer  visible,  and  bleach  the  iodine 

*  Gooch  and  Browning:  Methods  in  Analysis,  Gooch,  p.  291.  Gooch 
and  Morris :  Ibid.,  p.  294. 


192  QUANTITATIVE  CHEMICAL  ANALYSIS 

color  in  the  still  hot  liquid  by  the  cautious  addition  of  very  dilute  (ap- 
proximately N/100)  sulphurous  acid.  Should  too  much  sulphurous  acid 
be  inadvertently  added  the  error  may  be  corrected  by  allowing  the  hot 
solution  to  stand  in  the  air  until  the  iodine  color  returns,  when  the  bleach- 
ing may  be  repeated  with  more  care.  At  the  instant  when  the  bleaching 
has  been  accomplished  exactly,  dilute  the  solution  with  cold  water,  cool 
quickly,  and  nearly  neutralize  with  potassium  hydroxide.  Complete  the 
neutralization  with  potassium  hydrogen  carbonate,  and  add  an  excess  of 
this  reagent  amounting  to  about  20  cm.3  of  the  saturated  solution.  Titrate 
the  arsenite  in  the  solution  with  N/10  iodine,  using  the  starch  indicator. 

From  the  number  of  cubic  centimeters  (w)  of  iodine  used  and  the  iodine 
value  per  cm.3  (i),  calculate  the  weight  of  arsenic  (w)  from  the  proportion 
ni  \w  =  I2  :  As  =  2  X  126.92  :  74.96, 
ni  X  74.96 


w  = 


2  X  126.92 


Determination  of  Antimony  in  Antimonic  Acid.  —  In  treat- 
ing compounds  of  antimonic  acid  by  a  method  similar  to*  that 
just  described  for  the  determination  of  a  compound  of  arsenic 
acid,  it  is  necessary  to  take  the  precaution  of  preventing  the 
precipitation  of  antimonic  acid  from  solution  upon  the  ad- 
dition of  sulphuric  acid,  and  tartaric  acid  accomplishes  this 
effect  satisfactorily.  Furthermore,  the  separation  and  in- 
cipient volatilization  of  antimony  iodide  takes  place  at  a  mod- 
erate concentration.  Otherwise  the  process  employed  in  the 
estimation  of  arsenic,  arsenic  acid  and  its  derivatives  is  appli- 
cable to  the  determination  of  antimony  in  antimonic  acid  and 
its  derivatives,  according  to  the  reaction 

H3Sb04  +  2  HI  =  H3Sb03  +  H20  +  I2. 

According  to  procedure  developed  experimentally  an  antimo- 
nate,  in  amount  not  exceeding  the  equivalent  of  0.2  grm.  of 
antimony  trioxide,  is  treated  in  a  trapped  300-cm.3  Erlenmeyer 
flask  (see  Fig.  30)  with  tartaric  acid  (4  grm.)  and  sulphuric 
acid  to  acidity.  Thereafter  10  cm.3  of  sulphuric  acid  (1  :  1) 
and  1  grm.  of  potassium  iodide  are  added.  The  mixture  is 
boiled,  after  introducing  folded  platinum  foil  to  prevent  bump- 
ing of  the  liquid.  When  the  liquid  has  been  concentrated  to  a 
volume  of  about  50  cm.3  (not  less  than  45  cm.3)  the  boiling  is 
stopped,  and  the  color  of  the  iodine  is  bleached  by  the  cautious 
addition  of  sulphurous  acid  (approximately  N/100).  The  solu- 
tion is  diluted  at  once,  and  nearly  neutralized  with  potassium 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          193 

hydroxide.  The  neutralization  is  completed  with  potassium 
hydrogen  carbonate,  an  excess  of  this  reagent  (about  20  cm.3  of 
a  saturated  solution)  is  added,  and  the  titration  is  made  with 
N/10  iodine.  This  method  is  also  applicable  to  the  determi- 
nation of  antimony  and  arsenic  associated  in  the  higher  con- 
dition of  valence.*  It  may  also  be  used  in  the  determination 
of  either  of  these  elements  after  a  determination  of  associated 
copper  by  the  action  of  potassium  iodide  and  titration  of  the 
liberated  iodine,  according  to  the  method  described  on  p.  184, 
provided  that  citric  acid  is  the  acid  used  in  the  copper  separa- 
tion and  that  the  tetrathionate  formed  in  the  thiosulphate  titra- 
tion is  broken  up  by  treatment  with  bromine  before  proceeding 
further  with  the  method.f 

Evolution  of  Chlorine  and  Absorption  by  Potassium  Iodide: 
Titration  of  the  Iodine  set  Free 

Hydrochloric  acid,  which  is  not  susceptible  to  atmospheric 
oxidation  under  the  ordinary  conditions  of  action,  is  to  be  pre- 
ferred to  hydriodic  acid  wherever  it.  can  be  applied  in  processes 
in  which  the  halogen  evolved  in  the  reaction,  incomplete  in  the 
cold,  is  to  be  taken  as  the  measure  of  the  reducing  action  upon 
the  substance  to  be  determined.  In  such  processes,  the  evolved 
chlorine  is  made  to  react  with  potassium  iodide  and  sets  free  an 
equivalent  amount  of  iodine  which  thus  becomes  the  measure 
of  the  substance  to  be  determined. 

The  Operation.  —  The  operation  may  be  conducted  with  ad- 
vantage in  the  apparatus  shown  in  Fig.  29,  on  p.  187,  and  de- 
scribed in  the  context.  According  to  the  generally  applicable 
procedure,  the  weighed  material,  in  amount  not  exceeding  the 
equivalent  of  50  cm.3  of  N/10  thiosulphate,  is  put  into  the  re- 
action flask.  Connection  is  made,  as  shown,  with  the  trapped 
receiver  charged  with  a  solution  of  potassium  iodide  (3  grm.). 
The  stopper  is  inserted,  hydrochloric  acid  of  suitable  concen- 
tration is  introduced  through  the  stoppered  funnel,  a  current 
of  carbon  dioxide  is  sent  through  the  apparatus  (to  promote 
quiet  boiling  and  rapid  removal  of  the  liberated  chlorine),  and 
the  solution  in  the  flask  is  boiled  for  a  suitable  time.  The 
iodine  set  free  in  the  receiver,  ordinarily  titrated  with  N/10 

*  Gooch  and  Gruener:  Methods  in  Analysis,  Gooch,  pp.  308,  310. 
t  Heath:  Ibid.,  p.  318. 


194  QUANTITATIVE  CHEMICAL  ANALYSIS 

thiosulphate,  is  the  measure  of  the  substance  reacting  with  the 
hydrochloric  acid. 

Examples  of  this  procedure  are  determinations  of  manganese 
dioxide,  chlorates,  nitrates,  and  selenates. 

Determination  of  Manganese  Dioxide.  —  Suitable  propor- 
tions for  the  analysis  of  manganese  dioxide  are  0.2  grm.  of  this 
substance  with  50  cm.3  of  hydrochloric  acid  (1  :  2).  The  re- 
action in  the  flask  proceeds  according  to  the  equation 

Mn02  +  4  HC1  =  MnCl2  +  2  H20  +  C12, 
and  that  in  the  receiver  according  to  the  equation 
C12  +  2  KI  =  2  KC1  +  I2. 

In  this  process  the  chlorine  is  quickly  evolved  and  the  transfer 
to  the  receiver,  charged  with  potassium  iodide,  should  be  regu- 
lated with  care. 

Determination  of  Chlorates.  —  The  analysis  of  a  chlorate 
may  be  made  by  procedure  similar  to  that  employed  in  the 
treatment  of  manganese  dioxide.  In  consequence  of  the  high 
oxidizing  power  of  chlorates  the  amount  conveniently  handled 
is,  however,  comparatively  small.  With  potassium  chlorate, 
for  example,  the  reactions  are 

KC103  +  6  HC1  =     KC1  +  3  H2O  +  3  C12, 
3  C12  +  6  KI    =  6  KC1  +  3 12, 

and  the  amount  equivalent  to  50  cm.3  of  N/10  thiosulphate  is 
approximately  0.1  grm. 

Determination  of  Nitrates.  —  Manganese  chloride  in  con- 
centrated hydrochloric  acid  induces  the  easy  decomposition  of 
nitrates  with  the  production  of  nitrogen  dioxide  and  chlorine, 
appearing  itself  at  the  end  of  the  operation  in  its  original  form. 
At  the  ordinary  temperatures  the  action  is  slow,  but  upon 
warming  a  mixture  of  manganese  chloride,  hydrochloric  acid, 
and  a  nitrate,  the  decomposition  of  the  last  begins  with  the 
formation  of  a  high  chloride  of  manganese  and  liberation  of 
nitrogen  dioxide: 

2  HN03  +  3  MnCl2  +  6  HC1  =  3  MnCl4  +  4  H20  +  2  NO. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  195 

Ultimately,  if  the  heating  is  continued,  the  nitrogen  dioxide 
and  chlorine  from  the  higher  chloride  are  evolved  and  manga- 
nous  chloride  remains: 

3  MnCl4  =  3  MnCl2  +  3  C12. 

During  the  process  of  heating  the  color  of  the  solution  passes 
from  the  original  characteristic  green  of  manganous  chloride  in 
concentrated  hydrochloric  acid  through  darker  shades  to  black, 
and  returns  by  the  reverse  changes  to  the  original  tint.  In 
the  presence  of  concentrated  hydrochloric  acid  the  decompo- 
sition of  the  nitrate  extends  to  the  last  traces,  but  the  addition 
of  a  large  amount  of  water  will  reverse  an  action  already 
established  in  strong  acid.  The  reaction  serves  to  detect 
nitrates  (in  absence  of  other  oxidizers)  when  present  in  fairly 
small  amounts  (e.g.,  one  part  of  potassium  nitrate  in  sixty- 
thousand  parts  of  water)  and  may  be  utilized  in  the  quanti- 
tative estimation  of  nitrates.*  In  applying  the  reaction 
quantitatively,  the  nitrate  to  be  determined  is  treated  in  an 
atmosphere  of  carbon  dioxide,  with  a  saturated  solution  of 
crystallized  manganous  chloride  in  concentrated  hydrochloric 
acid,  the  volatile  products  of  action  —  chlorine,  nitrogen  dioxide, 
and  perhaps  nitrosyl  chloride  —  are  passed  into  a  solution  of 
potassium  iodide,  and  an  amount  of  iodine  equivalent  to  the 
nitrate  originally  present  is  set  free : 

3  C12  +  6  KI  =  6  KC1  +  3 12, 
or  6  NOC1  +  6  KI  =  6  KC1  +  6  NO  +  3 12. 

In  these  reactions  HNO3  o  3  I. 

The  apparatus  shown  in  Fig.  29  will  serve  for  the  operation 
if  a  trap  to  exclude  air  from  the  absorption  trap  and  receiver  is 
attached  outside  the  trap  sealed  to  the  receiver.  In  conducting 
the  experiment  the  receiver  and  traps  are  charged  with  a  solu- 
tion of  potassium  iodide  and  the  receiver  stands  in  a  jar  of 
cold  water.  The  nitrate  is  introduced  into  the  reaction  flask, 
and  followed  by  the  solution  (10  cm.3  for  every  0.1  grm.  of 
nitrate)  of  manganous  chloride  in  concentrated  hydrochloric 
acid.  The  current  of  carbon  dioxide  is  started  and,  after  a 
suitable  interval  for  the  removal  of  air,  heat  is  applied  and  the 

*  Gooch  and  Gruener:  Methods  in  Analysis,  Gooch,  p.  213. 


196  QUANTITATIVE  CHEMICAL  ANALYSIS 

distillation  is  continued  until  nearly  all  the  liquid  has  passed 
over.  The  iodine  in  the  receiver  and  trap  is  titrated  as  ex- 
peditiously  as  possible  by  N/10  thiosulphate  in  order  that 
traces  of  dissolved  nitrogen  dioxide  may  not  be  reoxidized  by 
the  air  and  again  react  to  liberate  more  iodine. 

Determination  of  Nitrites.  —  The  process  just  described  for 
the  determination  of  nitrates  is  also  applicable  to  the  determi- 
nation of  nitrites,  with  allowance,  of  course,  for  the  difference 
in  the  degree  of  oxidizing  power.  In  the  solution  of  manganous 
chloride  in  concentrated  hydrochloric  acid  the  nitrite  reacts 
according  to  the  equation 

2  HN02  +  MnCl2  +  2  HC1  =  MnCl4  +  2  H20  +  2  NO, 
the  manganese  chloride  breaks  up, 

MnCl4  =  MnCl2  +  C12, 
and  the  chlorine  sets  iodine  free  from  potassium  iodide, 

C12  +  2  KI  =  2  KC1  +  I2. 
In  these  reactions  HN02  ^  I. 

Evolution   of   Bromine   and   Absorption   by   Potassium 
Iodide.    Titration  of  the  Iodine  set  Free 

The  reaction  between  potassium  bromide,  a  strong  acid,  and 
certain  substances  capable  of  liberating  bromine  from  hydro- 
bromic  acid  has  been  put  to  advantageous  analytical  use. 
Examples  of  the  utility  of  the  reaction  are  processes  for  the 
determination  of  vanadates,  selenates,  and  tellurates. 

Determination  of  Vanadates.  —  The  reduction  of  vanadic 
acid  to  the  condition  represented  by  vanadium  tetroxide  by 
the  action  of  hydrochloric  acid  and  potassium  bromide*  is 
ideally  simple  and  complete  if  the  concentration  of  the  bromide 
is  kept  reasonably  small.  If,  however,  the  concentration  of 
hydrobromic  acid  (or  bromide)  is  too  great  the  reduction  is 
excessive  and  irregular.f 

In  the  practical  operation,  the  vanadate  (0.3  grm.  to  0.5 
grin.)  is  put  in  the  reduction  flask  of  the  distillation  apparatus 

*  Holverscheit:  Dissertation,  Berlin,  1890. 

t  Gooch  and  Curtis:  Methods  in  Analysis,  Gooch,  p.  337. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  197 

(see  Fig.  29)  with  potassium  bromide  (1.5  grm.  to  2  grm.), 
concentrated  hydrochloric  acid  (30  cm.3)  is  added,  and  the 
mixture  is  boiled  in  a  current  of  carbon  dioxide  until  the  color 
of  the  liquid  in  the  flask  is  a  pure  blue.  Bromine,  liberated 
according  to  the  reaction 

V2O5  +  2  HBr  =  V204  +  H20  +  Br2, 

is  absorbed  in  the  trapped  receiver  charged  with  potassium 
iodide  (3  grm.  in  300  cm.3)  and  the  liberated  iodine, 


=  2KBr  +  I2, 
is  titrated  with  N/10  sodium  thiosulphate. 

Determination  of  Selenates.  —  When  acted  upon  in  boiling 
solution  by  sulphuric  acid  and  potassium  bromide  in  suitable 
concentrations,*  selenic  acid  liberates  bromine  according  to  the 
reaction 

Se03  +  2  HBr  =  Se02  +  H20  +  Br2. 

When  the  concentrations  of  sulphuric  acid  and  the  bromide 
are  both  low  the  action  is  very  slow.  When  the  concentration 
of  the  bromide  is  large  bromine  is  retained  obstinately  in  the 
solution.  A  fairly  high  concentration  of  the  sulphuric  acid  and 
low  concentration  of  the  bromide  in  excess  make  the  conditions 
most  favorable  to  the  reaction.  Under  these  conditions  the 
bromine  is  removed  rapidly  to  the  distillate  leaving  a  colorless 
residue  which,  however,  again  takes  on  color  if  the  distillation 
is  prolonged,  while  elementary  selenium  makes  its  appearance 
in  the  distillate. 

In  handling  0.25  grm.  or  less  of  selenic  acid  in  an  initial 
volume  of  60  cm.3  containing  20  cm.3  of  sulphuric  acid  (1:1) 
and  1  grm.  of  potassium  bromide  the  reduction  is  almost  the- 
oretically exact  when  the  distillation  is  continued  until  the 
recoloration  of  the  boiling  liquid  is  barely  recognizable. 

The  apparatus  shown  in  Fig.  29  is  convenient  for  the  dis- 
tillation process.  The  bromine  evolved  is  absorbed  in  the 
solution  of  potassium  iodide  contained  in  the  receiver  and 
trap,  and  the  liberated  iodine  is  titrated  with  N/10  thiosulphate. 

*  Gooch  and  Scoville:  Methods  in  Analysis,  Gooch,  p.  386. 


198  QUANTITATIVE  CHEMICAL  ANALYSIS 

Determination  of  Tellurates.  —  By  the  action  of  potassium 
bromide  and  sulphuric  acid,  telluric  acid  may  be  reduced  to  the 
condition  of  tellurous  acid, 

Te03  +  2  HBr  =  Te02  +  H20  +  Br2, 

and  estimated  by  determining  the  iodine  liberated  from  potas- 
sium iodide  by  the  bromine  set  free  in  the  operation.* 

According  to  this  method,  the  alkali  tellurate  (equivalent  to 
0.5  grm.  of  tellurous  oxide,  or  less)  is  introduced  into  the  dis- 
tillation apparatus  (see  Fig.  29)  with  3  grm.  of  potassium 
bromide  and  10  cm.3  of  sulphuric  acid  (1  :  1)  in  a  total  volume 
of  50  cm.3  or  more.  A  current  of  carbon  dioxide  is  passed 
through  the  apparatus,  and  the  solution  is  boiled  to  set  free  the 
bromine,  which  is  absorbed  in  the  cooled  solution  of  potassium 
iodide  in  the  trapped  receiver.  The  iodine  liberated  from  the 
iodide  by  the  bromine  is  titrated  with  N/10  thiosulphate.* 

The  formation  of  tellurium  tetrabromide  in  the  concentrated 
acid  solution  makes  it  impossible  to  tell  by  the  color  when  all 
the  bromine  has  been  distilled,  but  the  experimental  evidence 
goes  to  show  that  the  boiling  of  the  liquid  from  a  volume  of  50 
cm.3  to  a  volume  of  25  cm.3  is  sufficient,  while  concentration 
from  100  cm.3  to  20  cm.3  apparently  does  no  harm. 

lodometric   Processes  Based  upon  a  Preliminary  Re- 
duction or  Oxidation  of  the  Substance  to  be 
Determined 

Certain  analytical  methods  demand  the  reduction  of  the  sub- 
stance to  be  determined  previous  to  a  titration  by  iodine.  The 
determination  may  rest  upon  the  action  of  standard  iodine  upon 
a  reduced  product  in  presence  of  an  excess  of  a  reducer  which  is 
inert  toward  iodine,  or  it  may  depend  upon  the  action  between 
iodine  and  the  remainder  of  a  known  amount  of  reducer  in  presence 
of  a  reduced  product  which  is  inert  toward  iodine. 

Determination  of  Vanadates:  Preliminary  Reduction  by 
Oxalic  Acid,  Tartaric  Acid,  or  Citric  Acid.  —  Vanadic  acid  may 
be  reduced  in  hot  solution  by  oxalic  acid,  tartaric  acid,  or 
citric  acid  and  reoxidized  by  iodine  in  alkaline  solution. f  To 

*  Gooch  and  Rowland:   Methods  in  Analysis,  Gooch,  p.  401. 
f  Browning  and  Goodman:  Ibid.,  p.  341. 


PROCEDURES  TN   VOLUMETRIC  ANALYSIS  199 

the  solution  of  a  vanadate  is  added  approximately  1  grm.  of 
the  organic  acid  for  every  0.1  grm.  of  vanadium  pentoxide  in  the 
substance  analyzed.  The  solution  is  heated  to  boiling.  To 
the  cooled  liquid  is  added  about  5  grm.  of  alkali  hydrogen  car- 
bonate for  every  gram  of  acid  used  and  iodine  in  slight  excess. 
The  mixture  is  set  aside  until  no  further  fading  of  the  iodine  is 
noticeable.  The  excess  of  iodine  is  bleached  with  standard 
arsenite,  starch  is  added,  and  the  excess  of  the  arsenite  is 
titrated  with  N/10  iodine.  The  reduction  by  oxalic  acid  takes 
place  in  the  sense  of  the  expression 

V2O5  +  H2C2O4  =  V2O4  +  H20  +  2  C02. 

Vanadium  tetroxide  is  also  a  product  of  the  reaction  with  tar- 
taric  or  citric  acids.  The  reoxidation  of  the  tetroxide  by  iodine 
in  the  presence  of  the  alkali  hydrogen  carbonate  used  to  take 
up  hydriodic  acid  takes  place  in  the  sense  of  the  expression 

V2O4  +  H2O  +  I2  =  V2O5  +  2  HI. 

From  the  iodine  value  per  cm.3  (i)  of  the  iodine  used  and  the  total  number 
of  cubic  centimeters  (n),  with  the  iodine  equivalent  per  cm.3  of  the  arsenite 
(i')  and  the  number  of  cubic  centimeters  (n'),  the  weight  of  vanadium 
pentoxide  (w)  may  be  found: 

=  1  grm.  mol.  V2O5  (ni  -  n'i')  =  182  (ni  -  rii') 
2  X  grm.  equivalent  I  253.84 

The  oxalic  acid  reduction  and  the  citric  acid  reduction  may 
be  applied  in  the  manner  described  to  the  determination  of  a 
vanadate  in  presence  of  a  molybdate  or  tungstate.  The  tar- 
taric  acid  process  may  be  used  to  determine  a  vanadate  in 
presence  of  a  tungstate  and  if  the  process  of  reduction  be  modi- 
fied by  substituting  digestion  in  the  cold  during  twenty-four 
hours  in  place  of  the  boiling  it  will  then  be  applicable  in  pres- 
ence of  a  molybdate  as  well  as  a  tungstate. 

Determination  of  Bromates.  Reduction  by  Arsenious 
Acid  with  Titration  of  the  Excess.  —  Bromic  acid  may  be 
reduced  by  arsenious  acid  in  acid  solution, 

3  H3As03  +  HBr03  =  3  H3AsO4  +  HBr, 


200  QUANTITATIVE  CHEMICAL  ANALYSIS 

and  determined  by  the  action  of  iodine  upon  the  arsenite 
formed  by  neutralization  of  the  residual  arsenious  acid.*  In 
this  process  it  is  sufficient  to  treat  the  bromate  with  a  con- 
siderable excess  of  N/10  arsenite,  acidify  into  sulphuric  acid 
[1  :  1]  (3  cm.3  to  7  cm.3),  make  up  the  solution  to  a  volume  not 
exceeding  200  cm.3,  boil  for  ten  minutes,  neutralize  with  alkali 
hydrogen  carbonate,  and  titrate  with  N/10  iodine. 

From  the  iodine  equivalent  per  cm.3  of  the  arsenite  (i)  and  the  number 
of  cubic  centimeters  employed  (ri),  with  the  iodine  value  per  cm.3  of  the 
iodine  (if),  and  the  number  of  cubic  centimeters  (nf),  the  amount  of  bromate 
is  given  by  the  proportion 

ni  -  rii'  :  w  =  6 1  :  RBr03, 

1  grm.  mol.  RBrO  3  (ni-rii') 
6  X  126.92 

Determination  of  Chromates:  Reduction  by  Arsenious 
Acid,  with  Titration  of  the  Excess.  —  Chromic  acid  is  reducible 
by  arsenious  acid,  in  presence  of  hydrochloric  acid,  in  accord- 
ance with  the  expression 

4  Cr03  +  3  AsaOs  +  n  As203  =  2  Cr203  +  3  Ast06  +  n  As203. 

According  to  procedure  based  upon  this  reaction  f  an  excess  of 
N/10  arsenite  is  added  to  the  cold  solution  of  a  chromate 
acidulated  with  10  cm.3  of  dilute  hydrochloric  acid  or  sulphuric 
acid  [1  :  3],  the  total  volume  being  less  than  100  cm.3.  After 
a  few  minutes,  when  the  solution  has  taken  on  the  bluish  green 
color  characteristic  of  chromic  salts,  the  solution  is  treated  with 
alkali  hydrogen  carbonate  in  excess  (about  5  grm.).  To  the 
alkaline  solution  N/10  iodine  is  added  in  excess,  the  mixture  is 
allowed  to  stand,  with  frequent  shaking,  for  about  half  an  hour, 
the  residual  iodine  is  taken  up  with  N/10  arsenite  and  the  ti- 
tration  is  finished  with  N/10  iodine  in  presence  of  starch. 

The  long  standing  is  necessary  on  account  of  the  tendency 
of  precipitated  chromic  hydroxide  to  hold  arsenious  acid  and 
thus  delay  the  oxidation  of  iodine.  The  presence  of  a  tart  rate 
before  the  neutralization  will  prevent  the  precipitation  but  the 

*  Gooch  and  Blake:   Methods  in  Analysis,  Gooch,  p.  474. 
t  Browning:  Ibid.,  p.  407. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          201 

dark  green  color  then  assumed  by  the  solution  makes  the  end 
reaction  of  the  starch  iodide  difficult  to  determine. 

In  the  presence  of  a  ferric  salt  the  determination  of  the  end 
reaction  is  also  difficult  in  consequence  of  the  formation  of  the 
reddish  brown  precipitate  unless  syrupy  phosphoric  acid  before 
the  neutralization,  in  which  case  the  iron  precipitate  will  be 
white.*  In  the  reactions  of  the  operation 

4  Cr03  =0=  3  As203  =0=  12 1. 

From  the  iodine  equivalent  per  cm.3  of  the  arsenite  (i)  and  the  number  of 
cubic  centimeters  (n),  with  the  iodine  value  per  cm.3  of  the  iodine  (i')  and 
the  number  of  cubic  centimeters  (n'),  the  weight  of  the  chromate  (w}  may 
be  calculated  from  the  proportion 

ni  -  n'i'  :  w  =  12  X  126.92:  4  grm.  mol.  M2CrO4, 
=  (ni  -  n'i')  (1  grm.  mol.  M2CrO4) 
3  X  126.92 

Determination  of  Tellurous  Oxide  by  Action  of  Perman- 
ganate and  Iodide,  with  Titration  of  the  Iodine  Set  Free.  — 
Tellurous  oxide  in  an  alkali  tellurite  may  be  determined  by 
treating  the  alkaline  solution  with  standard  permanganate  un- 
til the  pink  color  of  the  permanganate  or  the  green  color  of  the 
manganate  appears  in  the  meniscus  of  the  liquid,  adding  potas- 
sium iodide  with  sulphuric  acid,  and  titrating  the  iodine  set 
free.  The  free  iodine  may  be  titrated  by  standard  thiosulphatef 
or,  after  the  addition  of  an  excess  of  alkali  bicarbonate,  by 
standard  arsenite  (also  used  to  fix  the  standard  of  the  per- 
manganate) to  the  vanishing  point  of  the  iodine  color,  without 
starch.  J 

The  difference  between  the  iodine  found  and  the  iodine  equivalent  of  the 
permanganate,  Al,  measures  weight,  w,  of  the  tellurous  oxide  according  to 
the  proportion 

I2  :  TeO2  =  Al  :  w, 

=  Al  X  1  grm.  mol.  TeO2 
2  X  126.92 

*  Edgar:   Methods  in  Analysis,  Gooch,  p.  511. 

f  Norris  and  Fay:  Am.  Chem.  Jour.,  20,  278. 

t  Gooch  and  Peters:   Methods  in  Analysis,  Gooch,  p.  399. 


202  QUANTITATIVE  CHEMICAL  ANALYSIS 

Processes  Based  upon  the  Reaction  Between  Acids  and 
an  Iodide-iodate  Mixture 

The  action  of  certain  strong  acids  upon  the  neutral  solution 
containing  an  iodate  and  an  excess  of  a  metallic  iodide  results, 
under  suitable  conditions  of  concentration,  in  the  liberation  of 
iodine  in  definite  proportion  to  the  acid: 

MI03  +  5  MI  +  3  H2S04  =  3 12  +  3  M2S04  +  3  H20. 

The  action  is  non-reversible,  and  the  iodine  set  free  may  be 
determined  by  standard  sodium  thiosulphate  and  taken  as  the 
measure  of  the  acid.*  In  dilute  solutions,  however,  the  end- 
point  of  the  reaction  may  be  obscured  by  the  slow  and  con- 
tinuous evolution  of  iodine. f 

Determination  of  Hydrochloric  Acid,  and  Sulphuric  Acid.  - 
When  hydrochloric  acid  in  amounts  not  exceeding  0.15  grm.  or 
sulphuric  acid  not  exceeding  0.2  grm.  is  allowed  to  act  upon  a 
mixture  containing  1  grm.  of  potassium  iodide  and  0.1666  grm. 
of  potassium  iodate  in  a  total  volume  which  does  not  exceed 
100  cm.3  the  conditions  are  suitable  for  the  completion  of  the 
reaction  given  above,  and  the  liberated  iodine,  which  may  be 
titrated  by  sodium  thiosulphate,  is  an  exact  measure  of  the 
acid.t 

From  the  iodine  equivalent  per  cm.3  ({)  and  the  number  of  cubic  centi- 
meters (n)  of  the  thiosulphate,  the  weight  (w)  of  the  acid  may  be  calcu- 
lated: 

_  ni  X  1  grm.  mol.  HC1  _  ni  X  1  grm.  mol.  H2SO4 
126.92  2  X  126.92 

Boric  Acid.  —  Under  certain  conditions  of  dilution  at  which 
hydrochloric  acid  and  sulphuric  acid  will  react  completely  with 
the  iodide-iodate  mixture  boric  acid  has  no  appreciable  action 
upon  this  reagent.  The  iodide-iodate  mixture  may  be  used, 
therefore,  to  remove  free  hydrochloric  acid,  or  free  sulphuric 
acid,  from  a  solution  containing  free  boric  acid,  and  the  iodine 

*  Groger:  Zeit.  Angew.  Chem.,  1894,  52. 

t  Furry:  Am.  Chem.  Jour.,  6,  341. 

%  Phelps  and  Weed:  Methods  in  Analysis,  Gooch,  p.  60. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  203 

thus  liberated  may  be  taken  up  by  sodium  thiosulphate  while 
the  boric  acid  still  remains  free.* 

When,  however,  mannite  and  boric  acid  are  associated  the 
acidic  boromannite  compound  formed  is  sufficiently  strong  to 
liberate  from  the  iodide-iodate  mixture  an  amount  of  iodine 
equivalent  to  the  boric  acid  upon  the  assumption  that  the  acid 
reacts  in  the  monobasic  form,  as  metaboric  acid. 

5  KI  +  KI03  +  6  HOBO  =  3  I2  +  6  KOBO  +  3  H2O. 

Experimental  Process:  Determination  of  Boric  Acid.  —  In  the  practical 
utilization  of  these  reactions  for  the  determination  of  boric  acid  or  a 
borate,  take  in  a  solution  of  small  volume  an  amount  of  the  substance 
not  exceeding  the  equivalent  of  0.2  grm.  of  boric  anhydride,  add  hydro- 
chloric acid  (as  little  as  may  be),  shake  to  remove  any  excessive  amount 
of  carbon  dioxide,  and  nearly  neutralize  by  sodium  hydroxide  with  the 
use  of  litmus  paper  as  an  indicator.  Adjust  the  volume  of  the  still  dis- 
tinctly acid  solution  to  the  limit  of  50  cm.3,  and  treat  with  a  mixture  of 
potassium  iodide  (2  grm.)  and  potassium  iodate  (0.5  grm.).  Bleach  the 
iodine  liberated  by  the  free  hydrochloric  acid  with  a  strong  solution  of 
sodium  thiosulphate,  added  drop  by  drop  and  thoroughly  mixed  with  the 
solution,  and  add  N/10  iodine  to  faint  coloration.  Add  a  sufficient 
amount  of  mannite  to  saturate  the  solution  (10  grm.  to  15  grm.)  and  a 
measured  amount  of  N/5  thiosulphate  until  the  iodine  set  free  by  the 
boromannite  compound  is  bleached,  and  then  10  cm.3  more.  Again 
saturate  the  solution  with  mannite,  and  allow  it  to  stand  in  a  cool  place 
for  half  an  hour,  or  more.  Titrate  the  excess  of  thiosulphate  with  N/10 
iodine.  The  addition  of  the  starch  indicator  is  unnecessary,  since  half  a 
drop  of  the  N/10  iodine  is  sufficient  to  give  a  strong  coloration  to  solutions 
of  the  prescribed  volume. 

From  the  number  of  cubic  centimeters  (n)  and  the  iodine  equivalent 
per  cm.3  (i)  of  the  N/5  thiosulphate  solution,  with  the  number  of  cubic 
centimeters  (n')  and  the  iodine  value  per  cm.  (i'}  of  the  N/10  iodine,  cal- 
culate the  weight  (w)  of  boric  anhydride,  according  to  the  proportion 

ni  -  rii'  :  w  =  I2  :  B2O3  =  2  X  126.92  :  70, 

(ni  -  rii')  70 

2  X  126.92 

Determination  of  Acids  Liberated  in  the  Hydrolysis  of 
Salts.  —  Certain  salts  when  put  into  solution  in  water  tend  to 
undergo  hydrolysis  with  the  formation  of  a  basic  hydroxide,  or 

*  Jones:   Methods  in  Analysis,  Gooch,  p.  210. 


204  QUANTITATIVE  CHEMICAL  ANALYSIS 

basic  salt,  and  a  free  acid.  This  action  is  progressive  to  an  equi- 
librium and  is  generally  promoted  by  heat.  In  the  presence  of 
the  iodine-iodate  mixture,  the  free  acid  is  taken  up,  an  equiva- 
lent amount  of  iodine  being  set  free,  and  the  hydrolysis  advances 
to  completion  or  to  some  definite  stage  short  of  completion, 
depending  upon  the  nature  of  the  salt  and  the  concentrations. 
Aluminium  Sulphate  and  Aluminium  Chloride.  —  Though 
only  partially  hydrolyzed  at  ordinary  temperatures  in  presence 
of  the  iodide-iodate  mixture,  aluminium  sulphate  is  completely 
broken  up  by  heating,  according  to  the  reaction 

Ala  (S04)3  +  5  KI  +  KI03  +  3  H20  =  2  Al  (OH), 

+  3K2S04  +  3I2. 

The  behavior  of  aluminium  chloride  with  the  iodide-iodate 
mixture  is  similar.  Upon  the  presumption  that  the  alunu'nium 
salt  is  neutral,  this  reaction  affords  means  for  determining 
aluminium  iodometrically;  or,  it  may  be  applied  to  the  de- 
termination of  the  relation  of  the  acidic  ion  to  the  alumina  in 
the  salt  or  unknown  constitution  in  respect  to  chemical  neu- 
trality.* 

Other  Salts  Showing  Similar  Behavior.  —  Other  salts 
showing  complete  hydrolytic  decomposition  in  reaction  with 
the  iodide-iodate  mixture  are  chromic  sulphate  and  nickelous 
sulphate.  Ferrous  sulphate  and  cobalt  sulphate  are  similarly 
hydrolyzed  while  the  hydroxide  is  oxidized  at  the  expense  of 
the  excess  of  iodate.  The  hydrolysis  of  zinc  sulphate  comes 
to  an  end  with  the  formation  of  an  apparently  definite  basic 
sulphate. f 

Determination  of  Alkali  Hydroxides.  —  Alkali  hydroxides 
may  be  estimated  by  the  action  of  N/10  hydrochloric  acid  or 
N/10  sulphuric  acid  in  excess  and  the  determination  of  that 
excess  may  be  made  by  the  action  of  the  iodide-iodate  mixture 
(1  grm.  KI  :  0.166  grm.  KIO3;  50  cm.3  of  water)  and  titration 

*  Chemically  neutral  salts  of  alumina  undergo  partial  hydrolysis  in 
water  solution  and  show  an  acid  reaction  with  litmus  and  other  similar 
indicators. 

f  Moody:   Methods  in  Analysis,  Gooch,  p.  61. 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS          205 

of  the  liberated  iodine  with  N/10  thiosulphate  added  to  an 
initial  volume  not  exceeding  100  cm.3: 

MOH  +        (n  +  1)  HC1  =  MCI  +  H20      +  wHCl, 
6  nHCl  +  5  nKl  +  nKI03  =  3  n!2  +  6  nKCl  +  3  nH20. 

From  the  iodine  value  per  cm.3  (i)  and  the  number  of  centimeters  (n)  of 
the  thiosulphate  used,  with  the  iodine  equivalent  per  cm.3  (i'}  and  the 
number  of  cubic  centimeters  (n')  of  the  acid,  the  weight  (IP)  of  the  alkali 
hydroxide  will  be  given  by  the  proportion 

n'i'  -  ni  :  w  =  I  :  MOH, 

=  (ftV  -  ni)  1  grm.  mol.  MOH 
126.92 

Determination  of  Alkali  Carbonates.  —  Alkali  carbonates 
may  be  determined  by  the  action  of  N/10  sulphuric  acid 
in  excess,  boiling  to  remove  carbon  dioxide,  determination  of 
the  excess  of  acid  by  the  action  of  the  iodide-iodate  mixture 
(1  grm.  K  :  0.166  grm.  KIO3)  and  titration  of  the  free 
iodine  by  N/10  thiosulphate  in  an  initial  volume  not  exceeding 
100  cm.3, 

M2C03(1  +  3  n)H2S04  =  M2S04  +  3  nH2S04  +  H20  +  C02, 
5  nKl  +  nKIOa  +  3  nH2S04  =  3  nK2S04  +  3  nH20  +  3  nI2. 

From  the  iodine  equivalent  per  cm.3  (i)  of  the  thiosulphate  and  the 
number  of  cm.3  (n)  used,  with  the  iodine  equivalent  (i')  and  the  number 
of  cubic  centimeters  (n'}  of  the  acid,  the  weight  w  of  the  alkali  carbonate 
may  be  calculated  by 

n'i'  -  ni  :  w  =  2  I  :  M2CO3, 

_  (n'i'  —  ni)  1  grm.  mol.  M2COa 
2  X  126.92 

Processes  Dependent  upon  the  Action  of  Iodine  upon 
Basic  Hydroxides  to  form  an  Iodide  and 
an  lodate  in  Mixture 

Determination  of  Basic  Hydroxides.  —  When  the  solution 
of  an  alkali  hydroxide  is  acted  upon  by  iodine  at  a  temperature 
sufficiently  high  to  decompose  the  small  amount  of  hypoiodite 
that  might  otherwise  persist,  the  final  action  results  in  the 


206  QUANTITATIVE  CHEMICAL  ANALYSIS 

formation  of  an  exactly  neutral  mixture  of  iodide  and  iodate 
according  to  the  equation 

6  MOH  +  3  I2  =  MI03  +  5  MI  +  3  H2O. 
The  reaction  of  iodine  with  barium  hydroxide  is  similar: 
6  Ba(OH)2  +  6 12  =  Ba(lO3)2  +  5  BaI2  +  6  H2O. 

I.  If  a  known  amount  of  standard  iodine  is  used  to  bring 
about  the  reaction  under  such  conditions  that  the  excess  may 
be  kept  from  volatilization,  the  determination  of  the  excess  by 
standard  arsenite  will  give,  by  difference,  the  amount  used  up 
in  acting  upon  the  hydroxide,*  and  thus  supply  the  necessary 
data  for  determining  the  amount  of  the  hydroxide. 

The  operation  is  conveniently  carried  out  in  a  glass  stoppered 
flask  trapped  like  the  reaction-bottle  of  Fig.  28.  A  suitable 
measured  volume  (80  cm.3  to  90  cm.3)  of  N/10  iodine  is  •intro- 
duced into  the  flask  and  the  barium  hydroxide  is  added  either 
from  a  burette  or  from  a  stoppered  funnel  which  is  weighed  be- 
fore and  after  withdrawing  the  hydroxide.  The  glass  stopper  is 
introduced  and  the  trap  is  charged  with  a  solution  of  potassium 
iodide  to  prevent  the  escape  of  iodine.  The  solution  is  heated 
to  the  boiling  point  to  break  up  traces  of  hypoiodite  and  is  then 
cooled.  The  excess  of  iodine  is  determined  by  N/10  arsenite. 

From  the  iodine  value  per  cm.3  (i)  and  the  number  of  cubic  centimeters 
(n)  of  the  standard  iodine,  with  the  iodine  equivalent  per  cm.3  (if)  and  the 
number  of  cubic  centimeters  (n'}  of  the  standard  arsenite,  the  weight  (w) 
of  the  hydroxide  is  given  by  the  proportion 

ni  -  n'i'  :  w  =  I  :  MOH  =  21:  Ba(OH)2, 

_  (ni  —  n'i'}  1  grm.  mol.  MOH  _  (ni  —  n'i'}  1  grm.  mol.  Ba(OH)2 
126.92  2  X  126.92 

In  case  it  is  desired  to  record  the  iodine  value  per  cm.3  (i"}  of  the  hydrox- 
ide solution  (for  subsequent  use  in  the  differential  determination  of  an 
acid),f  this  value  may  be  found  by  taking  into  consideration  also  the 
number  of  cubic  centimeters  (n")  of  that  solution,  from  the  expression 

ni  —  n'i' 


II.    Instead  of  determining  the  amount  of  the  excess  of 
standard  iodine,  the  estimation  of  the  hydroxide  may  be  based 

*  Phelps:  Methods  in  Analysis,  Gooch,  p.  231.  t  See  p.  207. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          207 

upon  the  expulsion  of  the  excess  of  iodine,  the  treatment  of  the 
residue  of  iodide  and  iodate  with  sulphuric  acid  (or  hydro- 
chloric acid,  in  the  determination  of  barium  hydroxide)  and 
the  titration  of  the  liberated  iodine  by  N/10  thiosulphate.* 

From  the  iodine  equivalent  per  cm.  (i)  and  the  number  of  cubic  centi- 
meters (n)  of  the  thiosulphate  used  the  weight  w  of  the  hydroxide  is  given 
by  the  proportion 

ni:w  =  I:  MOH  =  I2  :  Ba(OH)2, 

_  ni  x  l  grm-  mo1-  MQH  =  ni  x  1  grm-  moL  Ba(OH)2 
126.92  2  X  126.92 

and  with  the  added  knowledge  of  the  number  of  cubic  centimeters  (n')  of 
the  hydroxide  solution  the  iodine  equivalent  per  cm.3  (i')  of  that  solution 
is  given  by  the  expression 


The  Differential  Determination  of  Acids.f  —  The  deter- 
mination of  acids  which  form  with  alkali  hydroxides,  or  with 
barium  hydroxide,  salts  which  are  not  susceptible  to  the  hydro- 
lytic  action  of  water,  may  be  determined  by  neutralizing  them 
with  the  solution  of  an  alkali  hydroxide,  or  the  solution  of 
barium  hydroxide,  of  previously  determined  iodine  value,  and 
then  determining  the  excess  of  the  hydroxide  remaining  (by 
one  of  the  methods  I  or  II  just  described). 

Hydrochloric  Acid.  —  Hydrochloric  acid  may  be  determined 
by  introducing  its  solution  into  an  Erlenmeyer  flask,  adding  a 
measured  excess  of  the  alkali  hydroxide  or  barium  hydroxide 
with  N/10  iodine  in  excess,  expelling  the  excess  of  iodine  by 
boiling,  cooling  the  residue,  and  then  acidulating  with  hydro- 
chloric acid  and  titrating  with  N/10  thiosulphate  the  iodine 
liberated  from  the  iodide  and  iodate. 

From  the  iodine  equivalent  per  cm.3  (i)  and  the  number  of  cubic  centi- 
meters (n)  of  the  thiosulphate  used,  with  the  iodine  equivalent  per  cm.3  (ir), 
and  the  number  of  cubic  centimeters  (nf)  of  the  standard  hydroxide,  the 
weight  w  of  the  acid  may  be  found  by  means  of  the  proportion 
rii'  -ni:w  =  I:  HC1, 

_  n'i'  -niXl  grm.  mol.  HC1 
126.92 

*  Walker  and  Gillespie:  Methods  in  Analysis,  Gooch,  p.  70. 
t  Ibid. 


208  QUANTITATIVE  CHEMICAL  ANALYSIS 

Sulphuric  Add.  —  Sulphuric  acid  may  be  determined  by  a 
process  precisely  similar  to  that  just  given  for  the  determination 
of  hydrochloric  acid. 

The  weight  (w)  will  be  found  by  the  expression 

_  n'i'  —  ni  X  1  grm.  mol.  H2SO4 
2  X  126.92 

Carbonic  Acid.  —  The  carbon  dioxide  of  carbonic  acid  may 
be  determined  by  absorbing  it  in  standardized  barium  hydrox- 
ide, thus  forming  the  insoluble  barium  carbonate,  and  estimat- 
ing the  excess  of  the  barium  hydroxide  from  the  excess  of  a 
measured  amount  of  standard  iodine  made  to  react  with  it 
according  to  the  process  previously  described  (see  p.  205). 

Processes  Based  upon  the  Action  of  Potassium  lodate 

in  Hydrochloric  Acid  upon  Substances 

Susceptible  to  Oxidation 

The  action  between  certain  substances  and  a  standard  iodate 
in  presence  of  hydrochloric  acid  (60  per  cent)  may  result  in  the 
reduction  of  a  part  of  the.iodic  acid  to  hydriodic  acid,  the 
interaction  of  these  acids  with  liberation  of  iodine,  and  the  con- 
version of  the  free  iodine  to  iodine  monochloride.  -Valuable 
processes  based  upon  this  characteristic  action  are  the  deter- 
minations of  arsenious  acid  and  antimonious  acid,* 

2  H3As03  +  KI03  +  2  HC1  =  2  H3As04  +  KC1  +  IC1  +  H20, 

KI03  o  2  H3As03, 
2  H3Sb03  +  KI03  +  2  HC1  =  2  H3Sb04  +  KC1  +  IC1  +  H20,  . 

KI03  <>2H3Sb03; 

processes  for  the  determination  of  sulphurous  acid  and  sul- 
phates, thiosulphates,  and  tetrathionates  (the  process  does  not 
apply  to  dithionates),f 
2  H2S03  +  KI03  +  2  HC1  =  2  H2SO4  +  IC1  +  KC1  +  H20, 

KIO3  =2H2S03  o2S02, 

Na2S203  +  2  KI03  +  2  HC1  =  Na2SO4  +  K2S04  +  2 IC1  +  H20, 
2  KI03  o  Na2S203, 

*  Andrews:  Jour.  Am.  Chem.  Soc.,  25,  756  (1908). 

t  Jamiesoh:  Am.  Jour.  Sci.  (4)  38,  166  (1914):  (4)  39,  639  (1915). 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS          209 

and  a  process  for  the  determination  of  copper  in  cuprous  sulpho- 
cyanide  which  is  adapted  to  the  estimation  of  copper  in  sul- 
phide ores.* 

(a)       Cu2(SCN)2  +  2  KI03  +  4  HC1  =  2  CuS04  +  3  KC1 

+  2  HCN  +  IC1  + 12  +  H2O, 

(6)       2 12  +  KI03  +  6  HC1  =  KC1  +  5 IC1  +  3  H20, 
or 

4  CuSCN  +  7  KI03  +  14  HC1  =  4  CuS04  +  7  KC1 

+  4HCN  +  7IC1  +  5HA 
KI03  *fCu. 

Experimental  Process:  Determination  of  Sulphurous  Add.  —  Introduce 
from  a  pipette  not  more  than  20  cm.3  of  a  carefully  measured  solution  of 
sulphurous  acid  below  the  surface  of  30  cm.3  of  cooled  concentrated  hydro- 
chloric acid  contained  in  a  250-cm.3  bottle  provided  with  a  glass  stopper. 
Add  chloroform  (6  cm.3)  and  measure  into  the  solution  from  a  burette 
standard  potassium  icdate  (about  ^V  mol.),  rapidly  and  with  gentle  mixing 
of  the  liquids,  until  the  iodine  set  free  at  first  has  largely  disappeared. 
Insert  the  stopper  in  the  bottle  and  shake  the  mixture  vigorously.  Con- 
tinue the  titration  gradually,  shaking  the  mixture  thoroughly  after  each 
addition  of  the  iodate,  until  the  violet  color  of  free  iodine  vanishes  from  the 
chloroform.  Calculate  the  amount  of  sulphuric  acid  from  the  expression 

KIO3  o  2  H2SO3 
derived  above. 

Experimental  Process  :f  Determination  of  Copper  in  Ore.  —  Treat  the 
copper  ore  (0.5  grm.)  in  a  200-cm.3  flask  with  a  strong  nitric  acid  (10  cm.3). 
Boil  off  the  larger  part  of  the  liquid  from  the  flask  kept  in  constant  motion 
and  inclined  to  prevent  mechanical  loss  in  the  process.  Cool,  add  concen- 
trated hydrochloric  acid  (unless  the  decomposition  of  the  ore  has  been 
already  accomplished)  and  boil  until  the  volume  of  liquid  amounts  to  about 
2  cm.3.  Cool,  add  12  cm.3  of  sulphuric  acid  (1:1),  and  boil  again  until 
fumes  of  sulphuric  acid  are  evolved  copiously.  After  cooling,  add  cold 
water  (25  cm.3),  heat  until  the  soluble  sulphates  dissolve,  filter,  and  wash 
the  precipitate  thoroughly.  Add  a  few  drops  of  hydrochloric  acid  and  if 
silver  chloride  is  precipitated  filter  again. 

Neutralize  the  filtered  solution  with  ammonium  hydroxide,  add  hydro- 
chloric acid  (not  exceeding  1  per  cent  by  volume  of  the  concentrated  acid) 
and  a  saturated  solution  of  sulphur  dioxide  (10  to  15  cm.3).  Heat  to  boil- 
ing and  add  ammonium  thiocyanate  (1  to  2  grm.)  in  solution,  stirring 

*  Wells  and  Jamieson:  Jour.  Am.  Chem.  Soc.,  30,  760*  (1908). 
t  Ibid. 


210  QUANTITATIVE  CHEMICAL  ANALYSIS 

thoroughly.     Settle  the  precipitate  (5  to  10  min.),  filter  on  paper  and  wash 
thoroughly  with  hot  water. 

Transfer  the  precipitate  and  paper  to  a  250-cm.3  bottle  provided  with  a 
glass  stopper,  add  20  cm.3  of  water,  30  cm.3  of  concentrated  hydrochloric 
acid,  and  6  cm.3  of  chloroform.  Introduce  standard  potassium  iodate  in 
successive  portions  from  a  burette,  inserting  the  stopper  and  shaking  the 
contents  of  the  bottle  vigorously  between  the  additions,  until  the  violet 
color  of  iodine  in  the  chloroform  just  disappears. 

Calculate  the  copper  from  the  expression 
KIO3  =0=  |  Cu 
derived  above. 

PRECIPITATION  PROCESSES 

Processes  may  frequently  be  made  available  in  which  a 
known  amount  of  a  precipitant  in  standard  solution  acts  upon 
a  chemically  equivalent  amount,  or  at  least  a  definite  amount, 
of  substance  in  solution  in  such  a  way  that  a  practically  in- 
soluble and  inert  product  is  formed.  In  precipitation  processes, 
the  end-point  may  be  fixed  either  by  noting  during  the  addition  of 
the  precipitant  the  point  at  which  no  further  precipitation  takes 
place,  the  point  at  which  a  precipitate  begins  to  form,  the  advent 
of  a  new  and  visible  reaction  in  the  solution,  or  the  first  indica- 
tion of  the  presence  of  an  excess  of  the  precipitant  in  a  series  of 
tests  made  upon  drops  removed  successively  from  the  solution. 

The  End-Point  Marked  by  Absence  of  Further 

Precipitation 

Determination  of  Silver  by  Precipitation  as  Silver  Chlo- 
ride. —  The  reaction  between  a  soluble  silver  salt  and  a 
soluble  chloride  forms  the  basis  of  a  classical  and  reliable  proc- 
ess for  the  exact  determination  of  silver  in  alloys  of  approxi- 
mately known  composition.  The  precipitant  is  sodium  chloride 
and  the  silver  is  precipitated  from  its  solution  in  nitric  acid, 
according  to  the  reaction 

NaCl  +  AgN03  <=±  AgCl  +  NaNa03. 

A  solution  of  pure  sodium  chloride  containing  2.7  grm.  in  1 
liter  (not  quite  the  equivalent  of  0.5  grm.  of  silver)  is  made  and 
from  this  a  second  solution  of  TV  concentration  is  made  by 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS          211 

suitably  diluting  a  portion  of  the  first.  The  stronger  solution 
is  standardized  against  pure  silver.  In  this  process,  an  amount 
of  silver  weighing  exactly  0.5  grm.  is  dissolved  by  heating  in 
a  (200-cm.3)  flask  with  10  cm.3  of  nitric  acid  (sp.  gr.  1.2),  the. 
free  oxides  of  nitrogen  are  blown  from  the  flask,  and  the  solution 
is  cooled.  After  adding  exactly  100  cm.3  of  the  stronger  salt 
solution  the  mixture  is  shaken  until  the  precipitate  is  coagu- 
lated and  settles  from  a  clear  supernatant  liquid.  To  this  clear 
liquid  is  added  from  a  burette  0.5  cm.3  of  the  sodium  chloride 
solution  of  7*0  concentration  so  that  the  liquid  will  run  down 
the  sides  of  the  flask  and,  spreading  over  the  surface  of  the 
liquid,  make  a  visible  cloud  of  silver  chloride.  The  liquid  is 
again  shaken  and  settled,  and  the  process  of  adding  0.5-cm.3 
portions  of  the  dilute  chloride  solution,  shaking,  and  settling  is 
repeated  until  the  addition  of  a  portion  of  0.5  cm.3  of  the  solu- 
tion causes  no  further  precipitation.  The  100  cm.3  of  the 
stronger  solution  with  the  measured  amount  of  the  weaker 
solution,  less  the  last  0.5  cm.3,  will  contain  the  amount  of  sodium 
chloride  which  corresponds  to  exactly  0.5  grm.  of  pure  silver. 
In  employing  the  solutions  thus  standardized  it  is  necessary, 
however,  in  order  that  the  highest  degree  of  accuracy  may  be 
attained,  to  have  present  in  the  analytical  process,  under  con- 
ditions otherwise  precisely  similar,  the  amount  of  silver  which 
was  present  in  the  standardization. 

Determination  of  Tellurous  Acid  by  Precipitation  of  Tel- 
lurous  Iodide.  —  Hydriodic  acid  and  tellurous  acid  interact 
with  the  formation  of  tellurium  tetraiodide,  converted  by 
water  into  an  oxyiodide  and  by  excess  of  alkali  iodides  to 
soluble  double  salts.  When  potassium  iodide  is  added  to  a 
cold  solution  of  tellurous  acid  containing  at  least  one-fourth 
of  its  volume  of  strong  sulphuric  acid,  no  tendency  to  form 
a  double  salt  becomes  apparent  until  the  potassium  iodide 
amounts  to  more  than  enough  to  convert  all  the  tellurous  acid 
present  into  tellurous  tetraiodide  according  to  the  equation 

H2Te03  +  4  H2S04  +  4  KI  <=>  TeI4  +  4  KHSO4  +  3  H2O. 

The  tellurium  tetraiodide  thus  formed  is  extremely  insoluble  in 
sulphuric  acid  of  one-fourth  strength,  though  soluble  in  excess 
of  potassium  iodide,  and  is  acted  upon  by  water  with  formation 
of  tellurium  oxyiodide  and  hydriodic  acid.  It  is  produced  at 


212  QUANTITATIVE  CHEMICAL  ANALYSIS 

first  in  the  condition  of  a  finely  divided  dark  brown  precipitate 
which  upon  agitation  of  the  liquid  gathers  in  curdy  masses  and 
settles,  leaving  the  liquid  clear.  By  taking  advantage  of  this 
tendency  to  curd,  it  is  possible  to  determine  without  great 
difficulty  the  exact  point  during  the  gradual  addition  of  potas- 
sium iodide  when  the  precipitation  of  the  tellurium  iodide  is 
complete.* 

Tellurous  oxide  is  dissolved  in  a  very  little  of  a  strong  solu- 
tion of  potassium  hydroxide,  and  dilute  sulphuric  acid  is  added 
carefully  until  the  tellurous  acid  which  is  precipitated  upon 
neutralization  of  the  alkali  hydroxide  is  just  redissolved.  To 
this  solution,  contained  in  an  Erlenmeyer  flask,  sulphuric  acid 
[1  :  1]  is  added  in  such  amount  that  the  liquid  shall  contain, 
after  the  subsequent  addition  of  potassium  iodide  in  solution, 
at  least  one-fourth  of  its  volume  of  strong  sulphuric « acid. 
The  flask  is  placed  upon  a  pane  of  window  glass  supported 
upon  strips  of  wood  about  1  cm.  above  the  level  of  the  table 
covered  with  white  paper.  A  solution  of  approximately  deci- 
normal  potassium  iodide,  free  from  iodate  and  carefully  stand- 
ardized,! is  introduced  gradually  from  a  burette  into  the 
middle  of  the  Erlenmeyer  beaker.  As  the  drops  of  the  potas- 
sium iodide  touch  the  liquid  the  precipitation  begins  at  the 
center  and  travels  in  rings  toward  the  outer  walls  of  the  beaker. 
When  the  liquid  becomes  so  opaque  that  a  newly  formed 
precipitate  is  distinguished  with  difficulty,  the  beaker  is  rotated 
and  the  curded  precipitate  permitted  to  settle;  and  then  the 
process  of  titration  is  continued  as  before  until  precipitation 
ceases.  In  an  Erlenmeyer  flask  10  cm.  in  diameter  across  the 
bottom  the  end  of  the  precipitation  is  easily  determined  if  the 
final  volume  of  the  liquid,  of  which  one-fourth  is  sulphuric  acid, 
does  not  exceed  100  cm.3.  In  this  process, 

Te02  o  4  KI 
according  to  the  equation  given  above. 

The  End-Point  Marked  by  the  Beginning  of 
Precipitation 

Determination  of  the  Cyanide  Ion.  —  When  a  solution  of 
silver  nitrate  is  added  to  an  excess  of  alkali  cyanide  the  pre- 

*  Gooch  and  Morgan:   Methods  in  Analysis,  Gooch,  p.  398. 
t  See  p.  177. 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS          213 

cipitate  of  silver  cyanide  first  formed  is  redissolved  by  the  ex- 
cess of  the  soluble  cyanide  with  the  formation  of  a  double  salt. 

AgNO3  +  KCN  <=>  AgCN  +  KN03, 

AgCN  +  KCN  <=>  AgK(CN)2. 

Upon  further  addition  of  silver  nitrate  the  soluble  double 
cyanide  is  attacked  and  silver  cyanide  is  precipitated. 

AgN03  +  AgK(CN)2  <=±  2  AgCN  +  KN03. 

In  applying  this  reaction  to  the  determination  of  the  cyanide 
ion  in  an  alkali  cyanide,  the  solution  is  suitably  diluted  before 
the  titration,  and  in  the  determination  of  hydrocyanic  acid  it  is 
first  treated  with  an  alkali  hydroxide  and  then  diluted.  The 
container  is  placed  upon  a  background  of  black  glazed  paper 
and  the  standard  silver  nitrate  is  added  cautiously,  and  with 
constant  stirring,  until  a  faint  turbidity  indicates  the  end-point 
of  the  reaction 

AgN03  +  2  KCN  <±  KAg(CN)2  +  KN03 
Ago2(CN). 

The  End-Point  Marked  by  a  New  Reaction 

Determination  of  the  Chloride  Ion  by  Precipitation  as  Silver 
Chloride,  with  an  Indicator  in  the  Solution.  —  In  determining 
silver  volumetrically  by  precipitation  as  silver  chloride  use  is 
made  of  the  reaction 

NaCl  +  AgN03  <=±  AgCl  +  NaN03, 

sodium  chloride  in  standard  solution  serving  as  the  precipitant 
for  silver  chloride  taken  in  the  form  of  the  nitrate.  In  making 
use  of  this  reaction  for  the  determination  of  the  chloride  ion, 
the  end-point  of  the  main  reaction  may  be  indicated  by  the 
advent  of  a  secondary  reaction,  provided  the  process  be  carried 
out  in  a  solution  which  is  either  neutral  or  devoid  of  free  acid 
excepting  acetic  acid,  the  first  excess  of  silver  nitrate  being  in- 
dicated by  its  action  upon  a  small  amount  of  potassium  chro- 
mate  (Mohr)  or  sodium  arsenate  (Lunge)  present  in  the  solution. 
Free  acids  other  than  acetic  acid  in  small  amount  are  eliminated 
by  the  addition  of  an  alkali  acetate,  and  solutions  of  colored 


214  QUANTITATIVE  CHEMICAL  ANALYSIS 

chlorides  of  the  metals  are  first  treated  with  an  alkali  hydroxide 
or  carbonate  to  remove  the  metal  hydroxide  and,  after  filtra- 
tion, with  acetic  acid  in  slight  excess. 

2  AgN03  +  K2Cr04  <=*  Ag2Cr04  +  2  KN03. 
2  AgN03  +  3  K2CrO4  +  2  HOC2H3O  <=»  Ag2CrO4  +  K2Cr207 

+  2  KN03  +  2  KOC2H3O  +  H20. 
3  AgN03  +  NaH2AsO4  +  2  NaOC2H3O  <=»  Ag3As04 
+  3  NaN03  +  2  HOC2H30. 

Determination  of  Silver  by  Precipitation  as  Silver  Thio- 
cyanate,  with  a  Ferric  Indicator  in  the  Solution.     ( Volhard.)  - 
The  titration  of  silver  nitrate  by  ammonium  thiocyanate  or 
potassium  thiocyanate, 

AgNO3  +  NEUSCN  ±5  AgSCN  +  NH4N03, 

with  the  aid  of  a  ferric  salt  to  indicate  the  point  when  the  thio- 
cyanate is  no  longer  engaged  in  the  reaction  of  precipitation, 
Fes  (S04)3  +  6  NH4SCN  <=>  2  Fe  (SCN)3  +  3  (NH4)2S04, 

is  another  process  illustrative  of  the  marking  of  the  end-point 
of  the  main  reaction  by  the  advent  of  a  new  reaction. 

Experimental  Process:   Precipitation  of  Silver  as  Silver  Thiocyanate. — 

A.  Preparation  of  the  Solutions  and  the  Indicator. 

1.  Prepare  an  approximately  N/10  solution  of  silver  nitrate  by 

dissolving  17  grm.  of  that  salt  in  water  and  making  the  solu- 
tion up  to  1  liter. 

2.  Prepare  an  approximately  N/10  solution  of  ammonium  thio- 

cyanate, NH4SCN,  by  dissolving  9  grm.  of  that  salt  in  water 
and  making  the  solution  up  to  1  liter. 

3.  Prepare  an  indicator  by  making  a  cold  saturated  solution  of 

ferric  alum  and  add  to  it  enough  of  pure,  colorless  nitric  acid 
to  bleach  the  brown  color  of  the  water  solution. 

B.  Standardization  of  the  Silver  Solution. 

Determine  exactly  the  standard  of  the  solution  of  silver  nitrate  by 
precipitating  the  silver  as  chloride  and  weighing  as  such,  ac- 
cording to  the  gravimetric  process. 

C.  Standardization  of  the  Thiocyanate  Solution. 

Put  a  measured  portion  of  the  standardized  silver  solution  (45  cm.3) 
in  a  beaker  with  an  equal  amount  of  water  and  1  cm.3  of  the 
ferric  indicator.  Add  a  measured  amount  of  the  thiocyanate 
solution  from  a  burette,  with  constant  stirring,  until  the  mixture 


PROCEDURES  IN  VOLUMETRIC  ANALYSIS  215 

of  precipitate  and  liquid  assumes  permanently  a  blush  of  red. 
Calculate  the  "silver  equivalent  per  cm.3"  of  the  thiocyanate 
solution. 
D.   Determination  of  Silver  in  a  Silver  Coin. 

Dissolve  a  carefully  weighed  ten-cent  piece  in  nitric  acid,  sp.  gr.  1.2; 
boil  the  solution,  dilute  with  water,  and  repeat  the  boiling  (to 
expel  nitrogen  oxides).  Cool  the  solution,  make  up  with  water 
to  a  suitable  volume  (200  cm.3)  in  a  graduated  flask.  Draw  from 
a  burette  a  measured  portion  of  this  solution  (40  cm.3),  add  the 
ferric  indicator  (1  cm.3),  and  titrate  as  in  the  standardization  of 
the  silver  solution.  Calculate  the  percentage  of  silver  in  the  coin. 

The  End-Point  Marked  by  Drop  Tests  Outside  the  Solution 
Determination  of  Phosphoric  Acid  by  Precipitation  as 
Uranyl  Phosphate,  with  a  Ferrocyanide  Indicator  Outside  the 
Solution.  (Pincus.)  —  When  the  solution  of  acid  potassium 
phosphate  is  treated  with  uranyl  acetate  a  greenish  white  pre- 
cipitate of  uranyl  phosphate  is  formed, 

KH2P04  +  U02  (C2H302)2  =  KC2H302  +  HC2H302  +  U02HP04. 

In  presence  of  ammonium  salts  the  precipitate  contains  am- 
monium uranyl  acetate,  U02NH4P04.  As  soon  as  the  uranyl 
acetate  is  in  excess  (in  the  boiling  solution)  a  drop  of  the  solu- 
tion touched  to  a  drop  of  potassium  ferrocyanide  upon  a  white 
surface  will  give  a  brown  coloration  which  marks  the  end-point. 
A  solution  of  uranyl  acetate  thus  standardized  against  potas- 
sium phosphate  or  calcium  phosphate  may  be  utilized  for  the 
titration  of  alkali  phosphate  or  calcium  phosphate,  respectively, 
under  similar  conditions  of  concentrations.* 

GASOMETRIC   PROCESSES 

Some  processes  of  analysis  in  which  definite  gaseous  products 
are  evolved  in  reaction  may  be  made  to  depend  upon  the  volu- 
metric measurement  of  the  gaseous  products.  Such  processes 
are  classed  under  the  heading  of  Gas-Volumetric  Analysis. 
Processes  in  which  gaseous  mixtures  are  analyzed  by  measure- 
ment of  the  original  substances  as  well  as  the  gaseous  products 
of  reaction  come  in  the  category  of  methods  of  Gas  Analysis. 

*  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed.,  Vol.  II,  p.  718. 


216 


QUANTITATIVE  CHEMICAL  ANALYSIS 


Gases  are  commonly  measured  over  a  liquid  in  some  form  of 
gas  burette  —  a  graduated  burette  provided  with  a  leveling  tube 
so  that  by  adjustment  of  the  liquid  in  the  burette  and  tube  to 
the  same  level  the  pressure  under  which 
the  gas  is  measured  may  be  equalized  with 
that  of  the  atmosphere.  (Fig.  31.) 

In  all  measurements  of  gases  it  is  neces- 
sary to  reduce  the  volume  as  measured, 
under  atmospheric  pressure  and  at  the 
atmospheric  temperature,  to  the  volume 
which  it  would  occupy  under  standard  con- 
ditions of  temperature  and  pressure,  0 
degree  and  760  mm.  Measured  in  dry 
condition  the  volume  (T^eo0)  of  the  gas  at 
0  degree  and  760  mm.  may  be  fourfd  from 
the  observed  volume  (V),  the  temperature 
(0,  and  the  barometric  height  (B)  cor- 
rected for  temperature  to  what  it  would 
be  at  0  degree  (fi)  by  means  of  the  formula 

l+# 


FIG.  31. 


in  which  a  represents  the  coefficient  of 
expansion  of  mercury  and  /5  the  linear 
coefficient  of  expansion  of  glass  —  the  cor- 
rection amounting  to  about  1,  2,  3,  4  mm.  for  temperatures 
lying  in  the  intervals  5  to  12,  13  to  20,  21  to  28,  29  to  35  degrees, 
respectively.  Then, 

VB0  VB0  (273) 

'  760  TAA  /1      i         1       j\  TAH  /"O7Q     I     -t\ 

7oU  (1  +  jf  j  t)        7oU  (276  +  t) 

If  the  measurement  is  made  over  water,  the  vapor  tension  of 
water  (w)  at  the  observed  temperature  must  be  taken  into  ac- 
count in  determining  the  pressure  (partial)  of  the  gas,  and  the 
expression  assumes  the  form 

0  =  V(B0-w)  273 
760    '  '    760(273  +  0 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS 


217 


When  many  reductions  of  gas  volumes  to  standard  conditions, 
are  to  be  made,  the  work  may  be  done  mechanically  with  the 
aid  of  a  device  in  which  the  gas  burette  and  leveling  tube  are 
connected  with  a  third  tube,  the  comparison  tube,  in  which 
air  is  confined,  the  amount  of  which  is  known  under  standard 
conditions.  Such  an  instrument  is  the  Lunge  Gas  Volumeter 
shown  in  Fig.  32.  This  instrument  consists  of  the  gas  burette 
(B)  (of  somewhat  more  than  100  cm.3  capacity), 
a  comparison  tube  (C),  and  the  leveling  tube 
(L),  all  connected  as  shown.  Air  is  introduced 
into  the  comparison  tube  (with  a  drop  of  water 
when  moist  gases  are  to  be  measured)  and 
adjusted  to  the  volume  which  corresponds 
under  the  prevailing  condition  of  temperature 
and  pressure  to  a  volume  of  100  cm.3  at  0  degree 
and  760  mm.,  as  determined  by  the  formula 
given  above.  That  is, 

760  (273  +  f) 


100 


(Bo  -  w)  273 


FIG.  32. 


In  measuring  a  gas  in  the  burette  (B)  the  tubes 

are  so  adjusted,  by  raising  or  lowering  them, 

that  the  level  in  the  comparison  tube  (C)  stands 

at  the  100  cm.3  mark,  while  the  levels  in  the 

burette  (B)  and  the  comparison  tube  (C)  are  the 

same.     Then  the  indicated  volume  in  the  burette  will  be  that 

of  the  contained  gas  at  0  degrees  and  760  mm. 

Gas-Volumetric  Methods 

The  determination  of  nitrates  by  measurement  of  the  nitro- 
gen dioxide  (NO)  which  they  are  capable  of  evolving,  and  the 
determination  of  ammonium  salts,  are  representative  processes 
of  gas-volumetric  analysis. 

The  Determination  of  Nitrates.  —  The  main  reaction  which 
takes  place  when  a  nitrate  is  acted  upon  by  hydrochloric  acid 
and  a  ferrous  salt  is  typified  by  the  equation 
2  KN03  +  8  HC1  +  6  FeCl2  =  2  KC1  +  6  FeCl3  +  4  H20  +  2  NO. 


218 


QUANTITATIVE  CHEMICAL  ANALYSIS 


The  reaction  is,  however,  accompanied  to  a  small  extent,  vary- 
ing with  the  conditions,  by  another  in  which  the  oxide  of  ni- 
trogen evolved  is  the  trioxide  N203,  but  the  higher  oxide  may 
be  reduced  to  the  lower  condition  by  contact  with  a  solution  of 
potassium  iodide: 

2  KI  +  H20  +  N203  =  2  KOH  +  I2  +  2  NO. 

The  results  most  nearly  in  accord  with  theory  of  the  main  re- 
action are  obtained  *  when  the  hydrochloric  acid  is  fairly  con- 
centrated, the  ferrous  salt  in  considerable  excess,  and  the  gas 
swept  from  the  boiling  liquid  through  a  solution  of  potassium 
iodide  to  the  measuring  burette  by  carbon  dioxide  and  col- 
lected over  sodium  hydroxide. 


FIG.  33. 

A  satisfactory  apparatus  for  the  operation  is  shown  in  Fig.  33. 
A  small  tubulated  retort,  upon  the  neck  of  which  is  fitted  a  small 
condenser,  serves  as  the  reaction  chamber.  Into  the  tubula- 
ture  of  this  retort  is  fitted,  by  a  ground  joint,  a  tube  drawn  out 
so  as  to  dip  below  the  surface  of  the  liquid,  and  fitted  to  stop- 
cocks, as  shown  in  the  figure,  and  so  branched  as  to  make  it 
possible  to  transmit  carbon  dioxide  through  the  apparatus,  or 
to  admit  a  liquid  without  introducing  air.  The  condenser  is 
joined  to  a  Will  and  Varrentrapp  bulb  containing  a  solution  of 
potassium  iodide,  and  this  is  connected  by  thick  vacuum  tub- 

*  Roberts:   Methods  in  Analysis,  Gooch,  p.  271. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  219 

ing  with  a  Hempel  gas  burette  charged  with  a  strong  solution 
of  sodium  hydroxide.  Carbon  dioxide  is  generated  in  a  Kipp's 
apparatus  by  action  of  boiled  hydrochloric  acid  (charged  with 
cuprous  chloride  to  take  up  dissolved  oxygen)  upon  boiled 
marble. 

In  using  this  apparatus,  the  nitrate  (about  0.1  grm.  of  potas- 
sium nitrate)  is  introduced  into  the  retort,  generally  in  the  dry 
condition,  carbon  dioxide  is  passed  through  the  apparatus  until 
the  gas  collected  over  sodium  hydroxide  leaves  only  the  minute 
bubble  which  the  gas  from  the  generator  alone  has  been  found  to 
give,  and,  after  shutting  off  the  carbon  dioxide  and  lowering  the 
leveling  tube  of  the  Hempel  burette,  a  boiled  solution  of  ferrous 
chloride  in  hydrochloric  acid  is  admitted  through  the  funnel 
tube.  The  liquid  is  then  slowly  heated  to  boiling  and  the 
process  continued  until  the  reaction  of  the  ferrous  salt  upon 
the  nitrate  is  apparently  complete,  when  the  carbon  dioxide  is 
again  passed  through  the  apparatus,  the  absorption  of  this 
gas  being  hastened  by  inclining  and  shaking  the  burette  at  in- 
tervals. The  volume  of  the  gas  residual  nitrogen  dioxide  (NO) 
is  noted  under  existing  barometric  and  thermometric  conditions 
and  the  volume  under  standard  conditions,  0  degree  and  760  mm., 
is  calculated  from  the  formula 

0         V  (B0-  w)  273 
760  (273  +  t)  ' 

Determination  of  Nitrogen  in  Ammonium  Salts.  —  When 
an  ammonium  salt  is  acted  upon  by  a  solution  containing  sodium 
hydroxide  and  sodium  hypobromite  the  reaction  proceeds  in 
the  sense  of  the  following  expression 

2  NH4C1  +  3  NaOBr  +  2  NaOH 

=  2  NaCl  +  2  NH4OH  +  3  NaOBr 
=  2  NaCl  +  3  NaBr  +  5  H2O  +  N2. 

At  100  degrees  the  reaction  might  be  completed  to  the  final 
stage  and  the  nitrogen  evolved  which  represents  the  entire 
amount  of  ammonium  salts,  but  at  the  atmospheric  temper- 
ature a  small  amount  of  the  ammonium  hypobromite  of  the 
first  stage  of  action  remains  undecomposed.  It  is,  however,  im- 
practicable to  apply  the  reaction  at  the  higher  temperature,  so 
that  it  is  best  to  find  empirically  the  relation  between  nitrogen 
evolved  and  the  ammonium  salt  from  which  it  is  evolved  under 


220  QUANTITATIVE  CHEMICAL  ANALYSIS 

given  conditions,  and  to  use  that  relation  instead  of  the  theo- 
retically equivalent  proportion  in  the  analysis  of  the  unknown 
ammonium  salt. 

The  nitrogen  may  be  collected  in  the  Lunge  apparatus, 
shown  in  Fig.  32,  the  adjustment  of  the  volume  to  standard 
condition   being   made   mechanically  as   pre- 
viously described.* 

Nitrogen  in  Organic  Substances.  —  When 
the  products  of  the  combustion  of  an  organic 
substance  mixed  with  copper  oxide  are  passed 
over  hot  copper  (to  decompose  oxides  of  nitro- 
gen), the  nitrogen  is  left  free,  and  if  the  opera- 
tion is  so  arranged  (Dumas)  that  air  is  excluded 
while  the  products  are  passed  on  (in  carbon 
dioxide)  to  the  azometer  (Fig.  34)  cnarged 
with  potassium  hydroxide,  the  amount  of  the 
nitrogen  may  be  determined  by  the  measure- 
ment of  its  volume. f 

Methods  of  Gas  Analysis 

The  practical  analysis  of  gas  mixtures  is 
"  34  generally  made  either  (I)  by  direct  treatment 

with  suitable  absorbents,  the  difference  between 
the  volumes  of  the  gases  before  and  after  treatment  measuring, 
the  amount  of  gas  absorbed,  or  (II)  by  first  submitting  the  gas 
mixture  to  process  of  chemical  treatment  (e.g.,  combustion) 
and  then  acting  upon  the  products  of  reaction  with  suitable 
absorbents. 

Many  forms  of  apparatus  (some  of  them  very  intricate)  find 
application  in  processes  of  gas  analysis.!  Of  these  the  com- 
bination of  the  "gas  burette"  and  "gas  pipette"  of  Hempel 
may  serve  as  a  simple  type.  The  "gas  burette"  consists  of 

*  P.  217.  Compare  also  Treadwell-Hall :  Analytical  Chemistry,  3rd 
ed.,  Vol.  II,  pp.  387,  822. 

t  For  details,  see  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed., 
Vol.  II,  p.  422. 

t  See  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed.,  Vol.  II,  p.  729, 
et  seq. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS 


221 


a  burette  (B)  and  a  leveling  tube  (L)  as  shown  in  Fig.  35. 
The  "gas  pipette,"  also  shown  (P),  holds  the  gas  mixture 
forced  in  from  the  measuring  tube  (by 
raising  the  leveling  tube)  for  the  action 
of  the  absorbent,  after  which  the 
residual  gas  is  returned  to  the  burette 
for  measurement.  In  case  it  is  neces- 
sary to  protect  the  absorbent  from  the 
action  of  air,  the  "double  pipette"  of 
Fig.  36  the  outer  bulbs  of  which  hold 
a  water  seal  is  serviceable. 

The  determination  of  the  oxygen 
content  of  air  may  serve  as  an  example 
of  the  simplest  of  the  methods  of  gas 
analysis. 

The  Determination  of  Oxygen,  and 
of  Nitrogen  with  the  Inert  Elements, 
in  Air.  —  The  oxygen  content  of  air 
may  be  determined  approximately  by 
direct  absorption  in  a  solution  of 
potassium  pyrogallate  contained  in  the 
double  pipette. 

In  the  preparation  of  the  pipette 
(Fig.  36),  enough  of  an  approximately  FIG.  35. 

25  per  cent  solution  of  pyrogallic  acid 

is  introduced,  from  a  funnel  attached  by  a  short  rubber  con- 
nector to  the  capillary  tube,  until  the  first  bulb  is  three-quarters 

full.  This  is  accomplished  by  applying 
alternately  gentle  suction  and  pres- 
sure by  blowing  at  the  outer  tube  of 
the  fourth  bulb  to  draw  in  liquid  or 
expel  air  as  may  be  necessary.  Next 
a  solution  of  potassium  hydroxide  (sp. 
gr.  1.27)  is  similarly  introduced  until 
the  liquid  fills  the  first  bulb  and  the 
FIG.  36.  connecting  tube  up  to  the  second  bulb. 

Then,  allowing  air  to  enter  the  capillary 

tube,  the  solution  is  drawn  into  the  second  bulb  until  it  is  two- 
thirds  full,  and  water  is  introduced  into  the  fourth  bulb.     By 


222  QUANTITATIVE  CHEMICAL  ANALYSIS 

gently  blowing  into  the  outlet  tube  the  water  is  made  to  flow 
into  the  third  bulb,  while  the  reagent  solution  fills  the  first 
bulb  and  the  capillary,  which  is  then  closed  with  a  cap.  The 
gradual  absorption  of  oxygen  from  the  air  inclosed  and  in  con- 
tact with  the  reagent  draws  in  more  water  from  the  fourth 
bulb.  In  the  final  adjustment  the  third  and  fourth  bulbs 
should  contain  enough  water  to  preserve  the  seal  when  the  air 
to  be  analyzed  is  forced  into  the  first  bulb.  With  the  first  bulb 
and  capillary  tube  filled  with  the  reagent  connection  is  made 
with  the  Hempel  burette.  Air  measured  under  atmospheric 
conditions  of  temperature  and  pressure  is  forced  from  the 
burette  into  the  pipette,  kept  in  contact  with  the  pyrogallate 
solution  (gently  agitated  with  a  rotary  motion)  for  three  or 
four  minutes,  returned  to  the  burette,  and  measured  again  under 
the  atmospheric  condition.  The  difference  between  the  volume 
of  the  air  taken  and  that  of  the  gas  remaining  (nitrogen  and  the 
inert  elements)  is  the  volume  of  oxygen  contained  in  air  meas- 
ured under  similar  conditions. 


COLORIMETRIC  PROCESSES 

The  fact  that  the  solution  of  a  colored  substance  shows  by 
transmitted  light  a  depth  of  color  proportionate  to  the  thickness  of 
the  layer  through  which  the  light  is  sent,  finds  application  in  most 
processes  of  colorimetric  analysis:  a  few  colorimetric  processes 
depend  upon  comparisons  of  colored  films  produced  under  similar 
conditions  of  reaction. 

Two  solutions  of  a  substance  in  the  same  solvent  will  have 
the  same  concentration  per  cm.3  when  equal  layers  having  the 
same  volume  transmit  the  same  depth  of  color.  When  varia- 
tion in  the  concentration  of  a  substance  in  solution  is  not  ac- 
companied by  chemical  change  (solvation,  hydrolysis,  ioniza- 
tion)  which  of  itself  has  effect  upon  the  color  of  the  solution, 
the  concentrations  per  cm.3  of  two  solutions  of  the  substance 
in  the  same  solvent  which  show  the  same  depth  of  color  for 
different  thicknesses  of  layers  through  which  the  light  is  trans- 
mitted, will  be  inversely  as  the  thicknesses  of  the  layers.  In 
such  cases,  the  concentration  per  cm.3  (cf)  of  an  unknown  solu- 
tion may  be  found,  in  terms  of  the  concentration  per  cm.3  (c) 
of  a  known  solution,  by  determining  the  thickness  of  the  layers 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  223 

(I1)  and  (I)  respectively,  which  transmit  the  same  depth  of 
color.     Then 

,      Ic 
C=T' 

If  variation  in  concentration  is  accompanied  by  chemical 
change  affecting  the  color,  then,  in  order  that  the  matching  of 
color  in  layers  of  equal  thickness  may  indicate  a  condition  of 
equal  concentrations,  it  is  necessary  that  the  layers  of  equal 
thickness  should  also  have  the  same  volume.  In  such  cases, 
the  solutions  compared  must  be  matched  in  color  by  adjusting 
the  volumes  (by  measured  dilution  or  strengthening  of  one 
solution  or  the  other)  until  the  equal  transmitting  layers  are 
also  equal  in  volume. 

Comparison  Tubes :  Colorimeters.  —  For  making  colori- 
metric  comparisons  many  different  forms  of  apparatus  have 
been  employed.  For  determinations  in  which  equal  volumes 
in  equal  layers  may  be  matched  in  color,  it  is  usual  to  employ 
sets  of  tubes  of  the  same  size  graduated  to  equal  volumes  — 
with  flat,  polished  bottoms,  for  lengthwise  observation  (such  as 
the  Nessler  cylinders  which  are  used  in  the  colorimetric  deter- 
mination of  small  amounts  of  ammonia  in  ammonium  salts) ; 
or  of  equal  thickness  of  wall  for  crosswise  observation  of  a 
clear  liquid  above  a  settled  precipitate  (such  as  the  tubes  used 
in  color  tests  for  carbon  in  steel).  For  most  purposes,  and 
especially  for  comparisons  of  colors  of  small  intensity,  an 
efficient  apparatus  is  a  simple  and  easily  constructed  colorim- 
eter *  consisting  of  two  graduated  comparison  tubes  of  equal 
size,  with  flat  and  polished  bottoms,  and  set  vertically  in  per- 
forations in  the  bottom  of  a  dark  box  (used  as  a  camera)  over 
a  mirror  which  is  illuminated  from  a  white  background.  If  a 
measured  amount  of  the  unknown  solution  is  put  into  one  of 
the  comparison  tubes  and  a  portion  of  the  known  solution  is 
measured  out  (best  from  a  burette)  into  the  other  tube  until 
the  colors  match,  the  amounts  contained  in  the  measured  por- 

*  For  other  excellent  devices  see  Hillebrand:  U.  S.  Geol.  Survey  Bulle- 
tin, 442,  pp.  33-38. 


224  QUANTITATIVE  CHEMICAL  ANALYSIS 

tions  of  the  known  and  unknown  solution  should  be  equal 
(provided  the  dilutions  are  suitable  in  respect  to  delicacy  of 
observation  and  that  differences  in  concentration  of  the  solu- 
tions have  no  chemical  effects).  Then,  dilution  of  the  liquids 
to  equal  volumes  will  not  change  the  matching  of  the  color. 
Should,  however,  any  appreciable  change  be  observed  on 
dilution,  the  test  may  be  regarded  as  preliminary  and  the  data 
used  to  point  the  way  to  making  up  the  solutions  (known  and 
unknown)  so  that  when  the  colors  of  equal  layers  match  the 
volumes  (and  concentrations)  will  be  practically  equal.  The 
amount  taken  of  the  substance  in  the  unknown  solution  will 
then  equal  the  amount  of  the  substance  in  the  known  solution 
measured. 

By  the  colorimeter  comparison  of  solutions  small  amounts  of 
manganese*  may  be  determined  in  the  form  of  permanganate; 
titanium  f  as  pertitanic  acid;  nitrous  acidj  as  tri-amido-azo- 
benzene  (Bismarck  brown)  or  as  indicated  by  a  mixture  of 
sulphanilic  acid  and  naphthylamine  in  acetic  acid;  and  gold  §  in 
red  colloidal  solution. 

The  Nephelometer.  —  Minute  amounts  of  material  in 
opalescent  suspension  can  be  determined  by  measuring  the 
brightness  of  light  reflected  on  the  suspended  particles  with 
reference  to  a  similar  suspension  of  known  content,  by  means 
of  the  Richards  Nephelometer.  \\  Two  test  tubes  slightly  in- 
clined toward  one  another  in  a  nearly  vertical  position  and 
arranged  so  as  to  be  partly  shielded  by  sliding  screens  from  the 
nearly  horizontal  rays  of  a  powerful  electric  light,  are  observed 
from  above  through  two  thin  prisms  which  bring  their  images 
together.  When  the  unknown  and  minute  quantity  of  a  dis- 
solved substance  (e.g.,  silver  chloride)  is  precipitated  as  a  faint 
opalescence  in  one  tube  and  a  known  amount  of  the  dissolved 
substance  is  treated  in  exactly  the  same  way  in  the  other  tube, 

*  Treadwell-Hall:  Analytical  Methods,  3rd  ed.,  Vol.  II,  p.  127. 

t  Ibid.,  pp.  100,  101. 

t  Ibid.,  p.  344. 

§  Maxsom:   Analytical  Methods,  Gooch,  p.  150. 

||  Richards  and  Wells:  Am.  Chem.  Jour.,  31,  235. 


PROCEDURES  IN   VOLUMETRIC  ANALYSIS  225 

the  light  reflected  from  the  one  tube  may  be  equalized  with 
that  reflected  from  the  other  tube  by  changing  the  position  of 
the  sliding  screens.  The  precipitates,  if  not  too  dense,  may  then 
be  taken  as  inversely  proportionate  to  the  lengths  exposed  to 
light.  Traces  of  substances  in  suspension  too  minute  to  be 
caught  by  any  ordinary  filter  (such  as  silver  chloride  in  opal- 
escent suspension)  may  be  thus  estimated. 


CHAPTER  V 

SYSTEMATIC   ANALYSIS 

Separations  and  Determinations 

The  Preparation  of  the  Substance.  —  Substances  in  gaseous 
condition  or  in  solution  are  generally  ready  for  the  immediate 
application  of  the  analytical  methods.  In  some  cases,  solid 
substances  may  also  be  submitted  at  once  to  the  determinative 
process,  as  in  the  determination  of  the  water  content^  of  a 
hydrous  salt  (see  p.  34)  or  in  the  determination  of  the  carbon 
dioxide  of  a  carbonate  (p.  38);  but  in  most  cases  the  first 
step  in  the  analysis  is  the  solution  or  decomposition  of  the  sub- 
stance by  means  of  a  liquid  reagent  adapted  to  the  character 
of  the  substance  and  the  nature  of  the  subsequent  operations. 
Common  reagents  for  this  purpose  are  water  (e.g.,  in  the  solu- 
tion of  copper  sulphate  —  see  p.  51),  nitric  acid  (e.g.,  in  the 
attack  upon  a  nickel-copper  coin  —  p.  53),  hydrochloric  acid 
(e.g.,  in  the  evolution  of  carbon  dioxide  from  calcite  —  p.  40; 
or  in  the  decomposition  of  hematite  —  p.  144),  aqua  regia  for 
gold  or  platinum,  or  hydrofluoric  acid  for  silica  and  certain 
silicates.  Sometimes  a  substance  which  is  refractory  to  acids 
or  ordinary  liquid  reagents  may  be  first  attacked  by  fusion  with 
a  suitable  flux  as  a  preliminary  to  processes  of  solution  —  e.g., 
the  fusion  of  silicates  with  sodium  carbonate  (see  p.  94),  and 
of  pyrites  with  sodium  carbonate  and  peroxide  (see  p.  78). 

It  should  not  be  forgotten  that  in  case  the  substance  contains 
organic  material  which  may  interfere  with  later  analytical  opera- 
tions —  e.g.,  tartaric  acid,  citric  acid,  or  other  organic  matter 
(sugar,  paraffin,  etc.)  in  large  amount  —  this,  in  general,  must 
be  removed  by  a  suitable  solvent  (water,  alcohol,  ether,  etc.) 
or  destroyed  by  ignition  or  by  the  action  of  an  oxidizing  agent 

226 


SYSTEMATIC  ANALYSIS  227 

(e.g.,  nitric  acid  or  the  mixture  of  that  acid  with  strong  sul- 
phuric acid)  before  proceeding  with  the  attempt  to  determine 
the  inorganic  constituents;  but  the  procedure  must  be  so  chosen 
that  it  will  not  occasion  the  loss  of  the  constituent  to  be  deter- 
mined or  affect  its  subsequent  determination  unfavorably. 

Qualitative  Tests.  —  In  most  cases  it  is  desirable  that  a 
thorough  qualitative  examination  should  precede  the  quanti- 
tative analysis,  at  least  so  far  as  to  make  it  sure  that  nothing 
is  present  which  may  interfere  with  the  accuracy  of  the  pro- 
posed procedure.  For  example,  the  analysis  of  a  preparation 
of  a  crystallized  hydrous  copper  sulphate,  for  the  purpose  of 
establishing  the  proportionate  composition  in  respect  to  cop- 
per, the  sulphate-ion,  and  water,  will  naturally  be  preceded  by 
a  qualitative  examination  at  least  sufficient  to  show  the  absence 
of  substances  which  if  present  might  interfere  with  the  analyti- 
cal process  —  e.g.,  the  electrolytic  determination  of  copper,  the 
precipitation  of  the  sulphate-ion  by  barium  chloride  and  weigh- 
ing as  barium  sulphate,  the  determination  of  water  by  loss  of  a 
portion  of  the  crystallized  substance  on  fusion  with  sodium 
paratungstate  (pp.  51,  76,  36). 

On  the  other  hand,  when  the  qualitative  analysis  is  to  be 
exhaustive  and  complete  so  far  as  may  be  practicable,  as  in 
the  more  elaborate  of  silicate  rocks  or  of  iron  ores,  the  prelimi- 
nary qualitative  examination  of  the  material  whose  compo- 
sition is  known  in  a  general  way  may  be  superfluous  inasmuch 
as  the  quantitative  methods  to  be  used  will  be  more  searching 
than  the  qualitative  procedure. 

The  Choice  of  Methods.  —  The  conditions  to  be  met  and 
the  difficulties  to  be  overcome  vary  as  widely  as  the  compo- 
sition of  substances.  The  determination  of  certain  selected 
constituents  of  a  substance  may  be  a  comparatively  simple 
matter  and  one  easily  attained  by  direct  and  special  methods. 
Such,  for  example,  are  determinations  of  iron  in  iron  ores  (p. 
144) ,  sulphur  in  pyrites  (p.  78) ,  or  phosphorus  in  iron  (p.  148) .  On 
the  other  hand,  the  complete  analysis  of  a  complex  substance 
may  prove  to  be  a  very  intricate  matter  involving  separations 
and  determinations  in  sequence  as  well  as  direct  methods  for 


228  QUANTITATIVE  CHEMICAL  ANALYSIS 

the  determination  of  certain  special  constituents.  In  the  deter- 
mination of  basic  ions  the  specifically  determinative  processes 
are  frequently  preceded  by  systematic  processes  of  separation 
which  involve  the  use  of  the  general  reagents  employed  in  qual- 
itative analysis  for  the  grouping  of  ions  of  reaction  according 
to  certain  characteristic  properties.  In  the  determination  of 
acidic  ions  group  reagents  (excepting  silver  nitrate  for  the  chlo- 
rine-ion, the  bromine-ion,  and  the  iodine-ion)  have  very  little 
application,  such  determinations  being  generally  made  by  direct 
and  special  methods,  which,  however,  often  involve  the  prelim- 
inary separation  of  interfering  substances. 

Grouping  of  Basic  Ions.  —  The  broad  classification  of  basic 
ions  according  to  their  behavior  in  respect  to  hydrogen  sulphide 
in  acid  solution,  to  an  alkali  sulphide,  and  to  ammonium 
hydroxide,  is  useful  in  quantitative  analyses,  as  in  qualitative 
analysis,  and  the  following  grouping  of  the  basic  ions  is  a  con- 
venient one. 

Group  I.  —  Basic  ions  whose  compounds,  soluble  in  water  or 
in  dilute  nitric  acid,  are  acted  upon  by  dilute  hydrochloric  acid 
with  the  formation  of  insoluble  chlorides.  (These  ions  may  also 
be  precipitated  as  sulphides  under  suitable  conditions.) 

Group  II.  —  Basic  ions  which  do  not  yield  chlorides  insoluble 
in  water  or  in  dilute  nitric  acid,  but  whose  compounds  soluble 
in  water  and  in  presence  of  dilute  hydrochloric  acid  are  acted 
upon  by  hydrogen  sulphide  with  the  formation  of  insoluble 
sulphides. 

Group  III.  —  Basic  ions  which  do  not  yield  insoluble  chlo- 
rides or  insoluble  sulphides  in  presence  of  dilute  hydrochloric 
acid,  but  whose  compounds  are  acted  upon  by  ammonium 
hydroxide  with  the  formation  of  hydroxides  (or  phosphates  if 
phosphoric  acid  be  present)  insoluble  in  a  faint  excess  of  the 
precipitant  and  in  presence  of  ammonium  salts,  and  whose 
acetates  are  hydrolyzed  when  heated  in  dilute  solution,  with  the 
formation  of  insoluble  basic  acetates  or  hydroxides. 

Group  IV.  —  Basic  ions  which  do  not  form  insoluble  chlo- 
rides or  sulphides  in  presence  of  dilute  hydrochloric  acid,  in- 
soluble hydroxides  in  presence  of  ammonium  salts,  or  insoluble 


SYSTEMATIC  ANALYSIS 


229 


basic  acetates,  but  whose  compounds  are  acted  upon  by  am- 
monium sulphide  with  the  formation  of  insoluble  sulphides. 

Group  V.  —  Basic  ions  which  do  not  yield  insoluble  chlorides 
or  sulphides  in  presence  of  dilute  hydrochloric  acid,  or  hydrox- 
ides in  presence  of  ammonium  salts,  or  basic  acetates,  but  whose 
compounds  are  acted  upon  by  ammonium  carbonate  in  presence 
of  ammonium  salts  in  dilute  water  solution  or  by  highly  con- 
centrated ammonium  carbonate  in  alcoholic  solution  (mag- 
nesium) with  the  formation  of  insoluble  carbonates. 

Typical  Forms  Precipitated  by  the  Group  Reagents  Applied  in  the  Order 

Indicated 


Group  I. 

Group  II. 

Group  III. 

Group  IV. 

Group  V. 

Group  VI. 

Precipitant 

NH4OH  in 

Precipitant 
HC1. 

Precipitant  H2S  in  acid 
solution,  after  the 
separation  of  Group  I. 

presence  of 
NH4C1, 
after  separa- 
tion of 
Groups  I 
and  II  and 

Precipitant 
(NH4)2  S, 
after  sepa- 
tion  of 
Groups 
I-III. 

Precipitant 
(NH4)2  C03, 
after  separa- 
tion of  Groups 
I-IV. 

Not  pre- 
cipitable  by 
the  group 
reagents  of 
Groups 
I-V. 

removal 

of  H2S. 

AgCl 

HgS      ] 

Al  (OH)3 

FeS 

BaC03 

Soluble 

Hg2Cl2 

PbS 

Fe  (OH)8 

NiS 

SrCO 

salts  of 

PbCl2* 

Bi3S3 

Cr  (OH)3 

CoS 

CaCOa 

Li 

T1C1* 

CuS 

Insol.  in 

Gl  (OH)2 

MnS 

f(NH4)2C03. 

Na 

CdS 

(NH4)2S 

+s 

Ti  (OH)4 
Zr  (OH)4 
Th  (OH)4 

ZnS 
(U02S) 

MgC03. 
(    4H20§ 

K 
Rb 
Cs 

Rh      Sn 
Pd 

Ce  (OH), 
Sc  (OH)  3 

NH4 

Os 

La  (OH), 

Partly 

Pr  (OH), 
Nd  (OH), 

PtS2 

sol.  in 

Sm  (OH)  3 

Au2So-3 

(NH4)2S 

+s 

Eu  (OH), 
Gd  (OH), 
Tb  (OH), 

Ir2S3 

Y  (OH), 

As2S3 

Dy  (OH), 

As2S5 

Ho  (OH), 

Sb2S3 
Sb2S5 
SnS 

Sol.  in 

(NH4)2S 

Er  (OH)3 
Tm  (OH), 
Yb  (OH), 
Lu  (OH), 

SnS2 

+s 

Ga  (OH), 
In  (OH), 

GeS 

Tl  (OH), 

Se'    6  2 

(NH4)2U2O7 
NH4V03| 

*  Precipitated  under  favorable  conditions  of  concentration,  always  incompletely. 

t  Precipitation  incomplete  in  presence  of  the  stronger  acids. 

t  In  a  solution  saturated  with  NH4CI. 

§  In  a  concentrated  alcoholic  solution  of  (NH4)2  CO3. 


230  QUANTITATIVE  CHEMICAL  ANALYSIS 

Group  VI.  —  Basic  ions  which  do  not  yield  insoluble  chlorides 
or  sulphides  in  presence  of  dilute  hydrochloric  acid,  insoluble 
hydroxides  precipitated  by  ammonium  hydroxide  in  presence 
of  ammonium  salts,  insoluble  sulphides  precipitated  by  am- 
monium sulphide,  or  insoluble  carbonates  precipitated  by  am- 
monium carbonate  in  aqueous  or  alcoholic  solution. 

These  groups  may  be  separated  by  the  application  of  the 
group  reagents  in  sequence  and  the  forms  of  combination  in 
which  they  appear  are  shown  in  the  preceding  table. 

Representative  Methods  of  Separation 

In  general,  the  separation  and  the  determination  of  both  basic 
ions  and  acidic  ions  may  be  accomplished  in  quantitative  anal- 
yses more  conveniently  by  the  application  of  special  methods 
rather  than  by  the  orderly  use  of  group  reagents  and  subsequent 
separations  of  the  members  of  a  group. 

The  Analysis  of  Brass 

Three  of  the  ordinary  constituents  of  brass  (copper,  with  lead 
and  tin  in  small  amounts)  are  precipitable  by  hydrogen  sulphide 
in  acid  solution,  one  (iron)  is  precipitable  by  ammonium  hydrox- 
ide, and  two  (zinc  and  traces  of  nickel)  are  precipitable  by 
ammonium  sulphide.  Instead,  however,  of  first  precipitating 
copper,  tin,  and  lead  together  from  a  suitably  prepared  solution 
and  afterward  separating  them,  it  is  preferable  to  remove  the 
tin  first,  by  special  treatment,  then  the  lead  by  another  special 
treatment,  and  then  the  copper  alone  by  means  of  the  hydrogen 
sulphide  in  acid  solution.  Similarly,  instead  of  precipitating 
the  zinc  and  nickel  together  by  ammonium  sulphide  and  then 
separating  them,  it  is  better  to  adjust  the  sulphide  precipitation 
in  such  a  way  that  the  zinc  only  may  be  first  precipitated  and 
the  nickel  recovered  from  the  filtrate  after  the  precipitation  of 
the  zinc. 

Experimental  Process:  Determination  of  Copper,  Zinc,  Tin,  Lead,  Iron, 
and  Nickel  in  Brass.  —  (1)  Solution  of  the  Metal.  Treat  in  porcelain  dish, 
covered  with  a  watch-glass,  a  portion  (0.5  grm.)  of  the  metal  with  cold 
nitric  acid  (sp.  gr.  1.2;  20  cm.3),  warming  as  the  violence  of  the  action 


SYSTEMATIC  ANALYSIS  231 

allows,  and  finally  evaporating  the  solution  to  dryness.  Moisten  the  dry 
residue  with  nitric  acid,  dissolve  the  soluble  portion  in  hot  water  (50  cm.3), 
filter  off  the  insoluble  residue  upon  ashless  paper,  washing  with  hot  water 
containing  ammonium  nitrate  (1  per  cent).  The  filtrate  contains  nearly 
all  of  the  copper,  zinc,  lead,  iron  and  nickel  as  nitrates.  The  residue 
is  chiefly  metastannic  acid  contaminated  with  traces  of  the  metals  other 
than  tin. 

(2)  Treatment  of  the  Residue  of  Impure  Metastannic  Acid.     Ignite  the 
precipitate  and  filter  in  porcelain  and  weigh  as  impure  stannic    oxide. 
Fuse  the  impure  oxide  with  a  mixture  of  sodium  carbonate  and  sulphur 
in  equal  parts  (six  times  the  weight  of  the  oxide)  in  the  covered  crucible 
until  the  flame  of  sulphur  which  burns  at  the  junction  of  cover  and  crucible 
is  no  longer  visible,  and  then  cool.     Treat  the  contents  of  the  crucible  with 
hot  water  and  should  any  gritty  residue  remain  separate  it  by  filtration, 
fuse  it  with  more  of  the  carbonate-sulphur  mixture,  and  again  treat  with 
hot  water.     To  the  solution  add  sodium  sulphite  until  brown  sodium 
polysulphide  is  changed  to  the  pale  yellow  monosulphide  and  filter  off  any 
precipitate  (metallic  sulphides  excepting  tin  sulphide)  and  dissolve  it  in 
nitric  acid,  adding  this  solution  to  the  filtrate  from  the  impure  metastannic 
acid.     To  the  sodium  sulphide  solution,  containing  the  tin  in  solution  as 
sulphostannate,  add  acetic  acid  to  acidity,  settle  the  precipitated  stannic 
sulphide,  filter,  wash  with  a  solution  of  ammonium  nitrate  (1  per  cent), 
ignite  the  precipitate  and  filter  (cautiously  at  first,  and  then  very  strongly) 
well  exposed  to  the  air,  cool,  add  a  little  ammonium  carbonate  (to  aid  in 
destroying  any  tin  sulphate)  and  reignite.     Weigh  the  residue  as  stannic 
oxide,  SnO2,  and  calculate  the  corresponding  amount  of  tin. 

(3)  Treatment  of  the  Filtrate  from  the  Impure  Metastannic  Acid  and  De- 
termination of  Lead  Sulphate.     Treat  the  filtrate  from  the  metastannic  acid, 
to  which  has  been  added  the  nitric  acid  solution  of  the  sulphides  left  in- 
soluble on  extracting  with  water  the  product  of  the  carbonate-sulphur 
fusion,  with  sulphuric  acid  (3  cm.3)  and  evaporate  on  the  water-bath  and 
over  the  radiator  until  all  nitric  acid  has  been  removed  and  the  dense 
fumes  of  sulphuric  acid  appear.     Add  water  (50  cm.3)  then  alcohol  (20 
cm.3),  stir  thoroughly,  filter  off  the  precipitate  upon  asbestos  in  the  per- 
forated crucible,  ignite  gently,  and  weigh  as  lead  sulphate,  PbSO4,  and 
calculate  the  corresponding  amount  of  lead. 

(4)  Treatment  of  the  Filtrate  from  Lead  Sulphate  and  Determination  of 
Copper.     Evaporate  the  filtrate  from  the  precipitated  lead  sulphate  to  re- 
move the  alcohol,  add  water  (100  cm.3),  heat  the  solution  to  boiling  and 
pass  in  hydrogen  sulphide  maintaining  the  current  while  the  liquid  cools. 
Filter  off  the  precipitated  cupric  sulphide,  washing  with  2  per  cent  sul- 
phuric acid.     Set  aside  the  filtrate  containing  the  iron,  zinc,  and  nickel  as 
sulphates.     Wash  the  precipitate  from  the  filter  as  well  as  may  be  into  a 
beaker  and  dissolve  it  with  nitric  acid.     Ignite  the  paper  and  adhering, 


232  QUANTITATIVE  CHEMICAL  ANALYSIS 

precipitate  and  dissolve  the  residual  copper  oxide  in  nitric  acid,  adding 
this  solution  to  the  solution  of  the  main  portion  of  the  precipitate.  Filter 
the  solution  of  copper  nitrate,  add  sulphuric  acid  (2  cm.3),  evaporate  to 
the  fuming  point  of  sulphuric  acid,  cool,  dilute  to  a  volume  of  50  cm.3 
and  electrolyze.  (See  pp.  49,  51.)  Weigh  the  deposit  of  metallic  copper. 

Instead  of  determining  the  copper  electrolytically,  it  may  be  precipitated, 
collected  on  asbestos,  and  weighed  as  cuprous  sulphocyanate*;  or  the  cupric 
sulphide  may  be  ignited  with  sulphur  in  hydrogen,  by  Rose's  method,  and 
weighed  as  cuprous  sulphide,  Cu2S.  (See  p.  105.) 

(5)  Treatment  of  the  Filtrate  from  Cupric  Sulphide  and  Determination  of 
Ferric  Oxide.     Boil  the  filtrate  from  the  precipitated  cupric  sulphide  to  re- 
move hydrogen  sulphide,  oxidize  the  ferrous  sulphide  by  means  of  bromine 
water,  and  add  ammonium  hydroxide  until  the  solution  smells  strongly  of 
ammonia.     Filter  off  and  wash  the  precipitated  ferric  hydroxide,  setting 
the  filtrate   aside.     Dissolve   the  precipitate  in   hydrochloric   acid   and 
reprecipitate  with  ammonium  hydroxide  (to  remove  any  included  zinc 
compound).     Filter  off  and  wash  the  precipitate,  adding  the  filtrate  to  the 
nitrate  from  the  former  precipitate  of  ferric  hydroxide.     Ignite  the  precipi- 
tate and  weigh  as  ferric  oxide,  Fe203.    Calculate  the  corresponding  amount 
of  iron. 

(6)  Treatment  of  the  Filtrates  from  Ferric  Hydroxide  and  Determination 
of  Zinc  Oxide.     Evaporate  the  nitrate  from  the  ferric  hydroxide,  acidify 
carefully  with  hydrochloric  acid  and  then  add  an  excess  of  the  acid  equiv- 
alent to  10  cm.3  of  the  N/10  acid  (0.0365  grm.)  to  every  100  cm.3  of  liquid. 
Heat  the  solution  to  60  or  70  degrees,  saturate  it  with  hydrogen  sulphide, 
and  set  it  aside  until  the  zinc  sulphide  has  settled.     Filter  off  the  precipi- 
tate, washing  with  a  2  per  cent  solution  of  ammonium  sulphate,  and  set 
the  filtrate  aside.     Dissolve  the  precipitate  in  hydrochloric  acid,  evapo- 
rate the  solution  (in  a  large  platinum  crucible)  to  remove  the  excess  of 
acid,  dissolve  the  residue  of  zinc  chloride  in  water,  add  mercuric  oxide  in 
large  excess,  evaporate  the  liquid,  ignite  the  residue,  and  weigh  as  zinc 
oxide,  ZnO.     Calculate  the  corresponding  amount  of  zinc. 

(7)  Treatment  of  the  Filtrate  from  Zinc  Sulphide  and  Determination  of 
Nickel.     Boil  the  filtrate  from  the  zinc  sulphide  to  expel  hydrogen  sulphide, 
add  to  the  hot  solution  a  1  per  cent  alcoholic  solution  of  dimethyl  glyoxime 
(CH3)2C2(NOH)2  (taking  care  that  the  alcohol  shall  make  only  a  small  pro- 
portion of  the  total  amount  of  liquid  —  never  more  than  half  —  to  avoid 
solubility  of  the  nickel  precipitate)  and  ammonium  hydroxide  in  faint  ex- 
cess.    Filter  on  asbestos  in  the  perforated  crucible,  wash  with  hot  water, 
dry   at    110   to    120    degrees,    and   weigh   as   nickel    dimethyl   glyoxime, 
(CH3)2C2(NO)2Ni  .  (CH3)2C2(NOH)2,  containing  20.31  per  cent  of  nickel. 
Calculate  the  amount  of  the  nickel. 

*  For  details,  see  Van  Name:  Methods  in  Analysis,  Gooch,  p.  108. 


SYSTEMATIC  ANALYSIS 


233 


(8)  State  the  results  of  the  analysis  in  the  following  form. 
ANALYSIS  OF  BRASS 


Weight  found. 

Weight 

taken  of 

brass. 

SnO2. 

PbS04. 

Cu. 

Fe203. 

ZnO. 

(CH3)2C2(NO)2Ni- 
(CH3)2C2  (NOH)2 

Grm. 

Grm. 

Grm. 

Grm. 

Grm. 

Grm. 

Grm. 

PERCENTAGE  COMPOSITION 

Copper Lead 

Zinc Iron 

Tin..  Nickel.. 


The  Analysis  of  Limestone 

The  chief  constituent  of  limestone,  calcium  carbonate,  is 
readily  attacked  by  hydrochloric  acid,  as  are  some  of  the  other 
constituents;  but  few  limestones  are  so  pure  that  this  treat- 
ment will  effect  complete  solution  or  even  complete  decompo- 
sition. The  residue  of  the  treatment  may  consist  of  quartz, 
silicates,  pyrites,  or  other  refractory  materials,  and  carbonaceous 
matter  may  also  be  found  in  the  residue.  The  thorough 
analysis  of  limestone  implies  the  use  of  a  very  elaborate  system 
of  procedures  and  the  determination  of  many  constituents,* 
but  for  many  purposes,  in  fact  for  most  technical  work,  the  ex- 
amination may  be  confined  to  the  determination  of  the  silica, 
alumina  (with  phosphoric  pentoxide  and  titanium  dioxide), 
ferric  oxide  (and,  perhaps,  ferrous  oxide),  calcium  oxide, 
magnesium  oxide,  carbon  dioxide,  and  water.  The  determi- 
nation of  these  substances  will  serve  as  another  illustration  of 
the  application  of  some  representative  procedures  in  systematic 
analysis. 

Experimental  Process:  Determination  of  CaO,  MgO,  FeO,  Fe20s,  Ak03  + 
P205,  Si02,  C02,  Insoluble  Siliceous  Matter,  and  Loss  on  Ignition.  —  Powder 

*  Hillebrand:  Bulletin  422,  U.  S.  Geol.  Survey,  p.  212. 


234  QUANTITATIVE  CHEMICAL  ANALYSIS 

the  material,  dry  at  100  degrees  (or  use  the  well  air-dried  material,  de- 
termining the  hygroscopic  water  in  a  small  portion  (1  grm.-5  grm.)  and 
calculate  determinations  to  the  basis  of  the  condition  of  the  material 
dried  at  100  degrees).  Weigh  out  exactly  for  analysis  several  portions 
A,  B,  C,  D,  E. 

A.  The  Determination  of  Loss  on  Ignition.     Weigh  a  portion  of  the 
material  in  a  weighed  platinum  crucible.     Ignite  to   constant  weight, 
gently  at  first,  and  finally  at  the  highest  heat  of  the  blowpipe.     Record  the 
loss  on  ignition  which  may  be  the  resultant  of  several  factors  —  loss  of 
water,  loss  of  carbon  dioxide,  possible  loss  by  oxidation  of  carbonaceous 
matter  to  volatile  carbon  dioxide,  and  possible  gain  in  weight  by  oxidation 
of  ferrous  iron  to  the  ferric  condition.     In  the  case  of  a  limestone  devoid 
of  carbonaceous  matter,  ferrous  compounds  and  sulphides,  the  loss  on 
ignition  will  probably  represent  fairly  well  the  sum  total  of  water  and 
carbon  dioxide.     By  subtracting  from  the  percentage  of  loss  on  ignition, 
the  percentage  of  carbon  dioxide  present  (and  to  be  calculated  from  the 
subsequent  determination,  B),   an   approximative  determination  of  the 
percentage  of  water  may  be  obtained. 

B.  The  Determination  of  Carbon  Dioxide.     Weigh  a  portion  of  the 
material  (1  grm.)  and  treat  it  with  hydrochloric  acid  in  an  apparatus  — 
e.g.,  the  Kreider  apparatus  —  suitable  for  the  determination  of  carbon 
dioxide  by  loss  (see  p.  40).     Calculate  the  percentage  of  carbon  dioxide. 

C.  The  Determination  of  Ferric  Oxide.     Put  a  weighed  portion  of  the 
material  (1  grm.)  in  a  small  beaker,  add  HC1  (10  cm.3),  cover  with  a  watch- 
glass,  digest  and  warm  until  the  acid  appears  to  have  no  further  action. 
When  the  ore  contains  organic  matter,  add  a  pinch  of  potassium  chlorate 
and  digest  until  effervescence  ceases.     Wash  off  the  watch-glass,  evaporate 
the  liquid,  treat  the  dry  residue  with  a  little  hydrochloric  acid  (5  cm.3), 
dilute  with  an  equal  amount  of  water,  filter  and  wash  upon  a  small  filter, 
allowing  the  liquid  to  run  into  a  small  flask  (50  cm.3  to  75  cm.3).     If  this 
treatment  of  the  ore  leaves  any  of  the  original  material  undecomposed, 
ignite  filter  and  residue  in  a  platinum  crucible,  treat  the  residue  with 
sulphuric  acid  (4  to  5  drops)  and  hydrofluoric  acid  (5  cm.3),  evaporate 
carefully  to  the  fuming  point  of  the  sulphuric  acid,  cool,  dilute  with  water, 
and  add  to  the  contents  of  flask.     Reduce  by  zinc  (see  p.  144)  the  ferric 
salt  in  solution,  hanging  a  funnel  in  the  neck  of  the  flask  to  prevent  loss 
by  spattering.     When  the  color  indicates  complete  reduction  add  15  cm.3 
of  sulphuric  acid.     When  all  the  zinc  is  dissolved,  pour  the  contents  of 
the  flask  into  a  liter  of  cold  water,  and  titrate  with  KMnO4  of  known 
standard.     From  this  amount  of  total  iron  thus  found,  calculate  the  cor- 
responding percentage  of  ferric  oxide.     In  case  no  ferrous  compound  is 
present  in  the  original  material  this  percentage  will  represent  the  ferric 
oxide.     If,  however,  a  compound  of  iron  in  the  ferrous  condition  is  sub- 
sequently found,  in  D,  the  percentage  of  ferric  oxide  corresponding  to 


SYSTEMATIC  ANALYSIS  235 

the  total  iron  must  be  diminished  by  that  corresponding  to  the  ferrous 
oxide  in  order  to  find  the  amount  originally  present  of  ferric  oxide. 

D.  The  Determination  of  Ferrous  Oxide.     Put  a  weighed  portion  of  the 
material  (1  grm.)  in  a  flask  (150  cm.3),  fill  the  flask  with  CO2,  add  JJC1 
(10  cm.3),  insert  a  two-hole  stopper  provided  with  inlet  and  outlet  holes, 
pass  a  current  of  CO2  through  the  flask,  and  warm  until  the  acid  appears 
to  have  no  further  action.     Cool  and  dilute  the  contents  of  the  flask  with 
cold,  freshly  boiled  water.     Filter  quickly  through  a  thin  felt  of  asbestos 
in  the  filtering  crucible,  wash  the  felt,  add  to  the  filtrate  and  washings 
15  cm.3  of  (1  :  1)  sulphuric  acid,  dilute  the  solution  with  boiled  and  cooled 
water  to  the  volume  of  1  liter,  and  titrate  with  potassium  permanganate. 

Remove  the  asbestos  and  residue  to  a  roomy  solid  platinum  crucible, 
washing  out  the  filtering  crucible  with  the  minimum  amount  of  water, 
add  H2SO4  (4  to  5  drops),  add  HF1  (5  cm.3),  fill  the  crucible  with  CO2,  put 
the  cover  in  place,  heat  the  contents  to  gentle  boiling  (in  order  to  keep 
the  crucible  full  of  steam),  immerse  crucible  and  contents  in  cold  water, 
add  5  grm.  of  calcium  chloride,*  1  grm.  of  manganous  chloride  (see  pp.  140) 
and  15  cm.3  of  sulphuric  acid  (1  :  1),  and  titrate  with  standard  KMnC>4. 

Add  together  the  percentages  of  ferrous  oxide  found  in  the  two  proc- 
esses of  titration  and  take  the  sum  as  indicating  the  ferrous  oxide. 

E.  The  Determination  of  Insoluble  Siliceous  Matter,  Silica,  Alumina, 
Calcium,  and  Magnesium.      (1)  Put  a  weighed  portion  of  the  material 
(1  grm.)  hi  a  small  dish  of  platinum,  fused  quartz,  or  (less  desirably) 
porcelain,  add  water  (25  cm.3),  cover  with  a  watch-glass,  and  then  cau- 
tiously add  hydrochloric  acid,  finally  warming  gently  to  complete  the 
action  of  the  acid  and  to  expel  dissolved  carbon  dioxide  while  the  dish  is 
still  covered.     Add  bromine  water  until  it  gives  its  color  to  the  solution 
(to  oxidize  ferrous  iron),  evaporate  the  solution  to  dryness,  and  desiccate 
at  110  degrees.     Treat  the  residue  with  hydrochloric  acid  (1  :  5),  warm 
and  digest  on  the  water-bath  for  some  minutes.     Filter  on  a  small  ashless 
filter,  wash  carefully  with  hot  water  and  preserve  the  filter  and  residue. 
Evaporate  the  filtrate  to  dryness,  desiccate,  extract  with  dilute  hydro- 
chloric acid,  filter,  wash  carefully  the  supplementary  residue  of  silica. 
Preserve  the  filtrate  for  further  treatment.     Ignite  together  in  platinum 
the    residues  of  both   evaporations,    and   weigh    as    Insoluble  Siliceous 
Matter. 

(2)  Fuse  the  insoluble  siliceous  matter  with  at  least  six  times  its  weight 
of  sodium  carbonate,  cool,  treat  the  mass  with  hydrochloric  acid  in  excess, 
evaporate  the  solution  to  dryness,  desiccate  at  110  degrees,  digest  the 
residue  with  hydrochloric  acid  (1  :  5),  filter,  and  wash  carefully  with  hot 
water.  Preserve  the  filter  with  the  precipitated  silica  upon  it. 

Evaporate  the  filtrate  from  the  precipitated  silica  to  dryness,  desiccate 
*  To  form  an  insoluble  fluoride  which  is  inert  toward  permanganate,  so 
that  the  end  reaction  may  be  definite. 


236  QUANTITATIVE  CHEMICAL  ANALYSIS 

at  110  degrees,  and  digest  the  residue  with  dilute  hydrochloric  acid  as  be- 
fore. Filter  off  the  supplementary  residue  of  silica  and  wash  carefully. 

Add  the  filtrate  to  the  filtrate  from  the  insoluble  siliceous  matter  of  (1). 

Ignite  together  the  first  residue  of  silica  and  the  supplementary  residue 
of  silica  obtained  in  the  evaporation  of  the  filtrate,  finishing  with  the  full 
heat  of  the  burner  or  blast-lamp.  Moisten  the  residue  with  sulphuric 
acid,  heat  gently  until  visible  fuming  ceases,  and  then  ignite  for  a  minute 
at  low  red  heat.  Weigh  as  impure  silica. 

Treat  the  impure  silica  with  hydrofluoric  acid  and  a  few  drops  of  sul- 
phuric acid,  evaporate  off  the  free  acid  over  a  radiator,  finally  ignite  to  low 
redness  for  a  minute,  and  weigh.  Take  the  loss  of  weight  of  the  impure 
silica  in  this  process  as  the  weight  of  silica. 

Fuse  with  sodium  carbonate  the  material  left  in  the  crucible  after  the 
treatment  with  hydrofluoric  acid  and  sulphuric  acid,  which  may  consist 
of  sodium  sulphate  with  sulphates  and  oxides  of  the  basic  elements  re- 
tained in  the  insoluble  siliceous  matter  when  the  limestone  was  acted  upon 
by  hydrochloric  acid,  and  add  the  solution  to  the  filtrate  from  the  insoluble 
.siliceous  matter  of  (1). 

(3)  To  the  filtrate  from  the  insoluble  siliceous  matter  of  (1),  to  which 
has  been  added  the  hydrochloric  solution  of  the  basic  elements  left  in  the 
hydrofluoric  acid  treatment  of  the  silica  of  (2),  add  hydrogen  dioxide 
(unless  manganese  is  absent)  and  ammonium  hydroxide  in  slight  excess. 
Heat  the  liquid  to  boiling,  filter,  and  wash  with  hot  water. 

Preserve  the  filtrate  for  further  treatment.  Dissolve  the  precipitate  off 
the  filter  in  hot  hydrochloric  acid  (1  :  1),  wash  the  filter  with  hot  water 
and  finally  with  dilute  ammonia.  Neutralize  the  acid  in  the  filtrate  by 
ammonium  hydroxide,  adding  a  slight  excess  of  the  last  (and  hydrogen 
dioxide  if  manganese  is  present),  filter  off  the  precipitate  upon  the  original 
filter,  and  wash.  Add  the  filtrate  to  the  former  filtrate  from  the  precipi- 
tate by  ammonium  hydroxide. 

Dry  the  precipitate  and  filter,  ignite  thoroughly,  and  weigh.  The 
residue  should  consist  of  ferric  oxide,  corresponding  to  the  total  iron  found 
in  C,  and  alumina  (with  manganese  oxide  (Mn3O4),  phosphoric  pentoxide 
(P206),  and  titanium  oxide  (TiO2),  in  case  manganese,  phosphorus  and  tita- 
nium were  constituents  of  the  limestone) .  The  difference  between  the  per- 
centage weight  of  this  residue  and  the  percentage  weight  of  ferric  oxide 
corresponding  to  the  total  iron  will  therefore  represent  the  percentage  of 
alumina  (with,  possibly,  oxides  of  manganese,  phosphorus,  and  titanium). 

(4)  To  the  combined  filtrates  from  the  precipitations  by  ammonium 
hydroxide  add  hydrochloric  acid  to  very  faint  acidity.     Heat  to  boiling 
and  add  slowly  a  solution  of  oxalic  acid  (1  grin.),  neutralize  the  acid  and 
finally  make  the  solution  ammoniacal  by  the  very  gradual  addition,  with 
constant  stirring,  of  1  per  cent  ammonium  hydroxide.     Digest  at  least  an 
hour  (over-night  if  the  amount  of  calcium  is  very  small).     Filter  and  wash, 


SYSTEMATIC  ANALYSIS 


237 


after  complete  settling,  and  set  the  nitrate  aside  for  further  treatment. 
Dissolve  the  precipitate  of  calcium  oxalate  (which  may  now  include 
magnesium  oxalate)  in  hot  (1:1)  hydrochloric  acid,  add  oxalic  acid,  make 
the  solution  ammoniacal,  digest,  filter  and  wash  this  second  precipitate  of 
calcium  oxalate  as  in  the  former  precipitation. 

Add  the  filtrate  to  the  filtrate  from  the  first  oxalate  precipitation. 

Dry  the  precipitate  and  paper,  ignite  (finally  with  the  blast-lamp),  and 
weigh  as  calcium  oxide, 

(5)  Evaporate  the  combined  filtrates  from  the  oxalate  precipitations  to 
a  volume  of  about  300  cm.3  and  neutralize  carefully  with  hydrochloric 
acid.     Add  hydrogen  sodium  ammonium  phosphate  (30  cm.3  of  the  satu- 
rated solution),  boil  for  five  minutes,  cool,  add  ammonium  hydroxide  until 
the  solution  smells  of  ammonia  acid,  then  add  10  cm.3  more  of  the  con- 
centrated ammonium  hydroxide.     Digest  until  the  precipitate  becomes 
crystalline.     Filter  off  the  precipitated  ammonium  magnesium  phosphate 
upon  asbestos  in  the  perforated  crucible,  heat  gently  until  ammonia  no 
longer  escapes,  ignite  at  bright  red  heat  and  weigh  as  magnesium  pyro- 
phosphate,   Mg^O?.     Calculate  the  corresponding  percentage  of  mag- 
nesium oxide. 

(6)  State  the  results  of  the  analysis  in  the  following  form. 

STATEMENT  OF  ANALYSIS 


Weight  taken, 


Weight  found 


SiO2 

A12O3,  etc 

Fe2O3 

FeO 

MgO.. 

CaO 

C02 

H2O  (approx.) 
Summation. . 


Grm. 


Per  cent 


The  Analysis  of  Silicates 

The  simple  analysis  of  a  silicate,*  involving  the  determination 
of  water,  silica,  aluminium,  iron,  calcium  and  magnesium,  follows 
very  closely  the  plan  of  the  analysis  of  limestone,  as  described 
above,  provided  the  silicate  is  decomposable  by  hydrochloric 
acid.  Generally,  however,  the  decomposition  of  the  silicate 

*  For  full  information  in  respect  to  the  analysis  of  silicate  rocks,  see 
Hillebrand:  Bulletin  422  of  the  U.  S.  Geological  Survey. 


238  QUANTITATIVE  CHEMICAL  ANALYSIS 

is  best  accomplished  by  fusion  with  sodium  carbonate.  The 
melt  is  then  cooled,  decomposed  by  hydrochloric  acid,  and  sub- 
jected to  the  processes  described  for  the  successive  separations 
and  determinations  in  the  analysis  of  the  insoluble  siliceous 
residue  of  the  limestone. 

The  determinations  of  the  alkali  elements  and  of  fluorine  are 
made  by  special  treatment  of  separate  portions  of  the  material 
for  analysis. 

Determination  of  Alkali  Elements.  —  The  alkali  elements, 
not  precipitable  by  ordinary  group  reagents,  may  be  converted 
to  weighable  form  and  determined  after  the  separation  of  other 
basic  elements.  Potassium  and  sodium  may  be  conveniently 
weighed  as  chloride,  sulphate  or  pyrosulphate.*  From  a 
mixture  of  potassium  chloride  and  sodium  chloride,  potassium 
may  be  precipitated  as  potassium  chlorplatinate  or  as  potassium 
perchlorate  (both  insoluble  in  alcohol)  and  determined  gravi- 
metrically,  or  the  potassium  may  be  determined  volumetrically 
by  a  process  involving  the  precipitation  of  potassium  sodium 
cobaltinitrite,  K2NaCo(N02)6  •  H20,  and  the  oxidation  of  this 
product  by  standard  potassium  permanganate,  this  process 
being  especially  applicable  to  the  determination  of  potassium 
in  soils,  fertilizers,  and  animal  fluids.! 

When  the  weight  of  a  mixture  of  potassium  chloride  and  so- 
dium chloride  is  known,  the  difference  between  that  weight  and 
the  weight  of  potassium  chloride  found  experimentally,  by  one 
of  the  methods  mentioned,  measures  the  sodium  chloride.  If 
the  potassium  is  removed  from  such  a  mixture  by  precipitation 
as  potassium  perchlorate,  by  means  of  perchloric  acid{  or 
aniline  perchlorate  §  in  alcoholic  solution  (97  per  cent),  the 
sodium  may  be  precipitated  as  chloride  by  saturating  the  alco- 
holic filtrate  (97  per  cent)  with  gaseous  hydrogen  chloride,  and 
the  sodium  chloride  thus  precipitated  may  be  determined  by 

*  Browning:  Methods  in  Analysis,  Gooch,  p.  79. 
f  Drushel:  Methods  in  Analysis,  Gooch,  p.  93. 

t  Treadwell-Hall:    Analytical    Chemistry,    3rd     ed.,    Vol.    II,    p.    50. 
Kreider:   Methods  in  Analysis,  Gooch,  p.  88. 
§  D.  U.  Hill:  Am.  Jour.  Sci.,  40,  75  (1915). 


SYSTEMATIC  ANALYSIS  239 

collecting  it  on  asbestos  in  the  perforated  crucible,  washing 
with  alcohol  (97  per  cent),  drying  at  about  110  degrees  and 
weighing  as  such.* 

The  determination  of  potassium  and  sodium  in  silicates  in- 
volves first  of  all  the  decomposition  of  the  silicates,  and  this 
may  be  brought  about  according  to  the  nature  of  the  case,  by 
the  action  of  hydrochloric  acid,  by  hydrofluoric  acid  and  sul- 
phuric acid  (Berzelius),  or  by  ignition  in  mixture  with  calcium 
carbonate  and  ammonium  chloride  (J.  Lawrence  Smith).  The 
last  method  is  of  wide  application  and  very  direct  in  consequence 
of  the  fact  that  when  the  residue  of  the  ignition  is  extracted 
with  water  the  silica  and  all  basic  elements  other  than  the  alkali 
elements  and  calcium  are  precipitated  by  the  calcium  hydroxide 
produced  in  the  process.  The  determination  of  potassium  and 
sodium  in  feldspar  will  illustrate  the  method.f 

Experimental  Process:  Determination  of  Potassium  and  Sodium  in 
Feldspar.  —  Process  of  Decomposition.  Weigh  out  a  portion  of  the  finely 
ground  material  (0.5  grm.),  put  it  into  a  roomy  platinum  crucible,  mix  it 
carefully  with  an  equal  weight  of  pure  sublimed  ammonium  chloride,  add 
eight  times  its  weight  (4  grm.)  of  calcium  carbonate  (prepared  by  precipi- 
tation with  ammonium  carbonate  and  of  determined  purity),  and  mix  the 
whole  with  great  care.  Heat  the  covered  crucible  over  a  low  flame  so 
regulated  that  the  ammonium  chloride  will  be  decomposed  without  vola- 
tilizing as  such,  while  calcium  chloride  (the  active  agent  in  the  attack  upon 
the  silicate)  is  formed  and  ammonia  is  set  free.  When  the  evolution  of  am- 
monia has  ceased,  raise  the  temperature  and  ignite  for  an  hour  in  such 
fashion  that  only  the  lower  half  of  the  crucible  (conveniently  held  in  an 
ignition-plate  J)  is  heated  to  low  redness.  Cool,  transfer  the  fritted  mass 
(using  water  if  necessary)  to  a  dish  (preferably  platinum),  digest  with 
water  (75  cm.3)  kept  hot,  pulverize  with  a  pestle  any  large  particles,  settle 
the  insoluble  residue,  decant  upon  a  filter  the  clear  solution  now  containing 
chlorides  of  the  alkali  elements  and  calcium,  and  wash  the  residue  thor- 
oughly by  decantation  and  afterwards  upon  the  filter. 

Removal  of  Calcium.  —  To  the  filtrate,  heated  to  boiling,  add  ammonium 
carbonate  and  ammonium  hydroxide,  filter  the  solution  of  alkali  chlorides 
from  the  precipitated  calcium  carbonate  and  set  the  filtrate  aside.  Dis- 

*  D.  U.  Hill.    Am.  Jour.  Sci.,  40,  75  (1915). 

f  For  the  determination  of  lithium  associated  with  potassium  and 
sodium,  see  Treadwell-Hall :  Analytical  Chemistry,  3rd  ed.,  Vol.  II,  p.  53. 

t  See  footnote  on  p.  78. 


240  QUANTITATIVE  CHEMICAL  ANALYSIS 

solve  the  precipitate  in  hydrochloric  acid  and  reprecipitate  with  ammo- 
nium carbonate  and  ammonium  hydroxide,  filtering  and  washing  to  recover 
any  alkali  chloride  which  may  have  been  included  in  the  former  precipita- 
tion. Evaporate  the  combined  filtrates,  dry  the  residue  thoroughly  at 
110  degrees  and  remove  ammonium  chloride  by  cautious  ignition.  Dis- 
solve the  residue  in  a  little  water,  and  remove  residual  traces  of  calcium 
by  treatment  of  the  solution  with  ammonium  oxalate  and  ammonium 
hydroxide,  standing  over  night,  filtering,  and  washing. 

Recovery  of  Alkali  Chlorides.  —  Evaporate  the  filtrate  in  a  dish  or  large 
crucible  of  platinum,  dry  the  residue  carefully  and  ignite  gently  to  remove 
the  ammonium  salts.  Moisten  the  residue  with  hydrochloric  acid  (to 
recover  as  chloride  any  alkali  salt  changed  to  carbonate  in  the  ignition), 
ignite  gently  and  weigh  the  alkali  chlorides. 

Separation  and  Determination  of  Potassium.  —  Dissolve  the  alkali 
chlorides  in  a  little  water,  add  in  water-solution  a  moderate  (calculated) 
excess  of  chlorplatinic  acid,  H2PtCl6,  and  evaporate  the  solution  on  the 
water-bath  until  the  residue  will  solidify  when  allowed  to  cool.  To  the 
cooled  residue  add  a  little  absolute  alcohol,  disintegrate  the  solid  mass  by 
means  of  a  stirring  rod  with  rounded  end,  filter  off  the  precipitate  upon 
asbestos  in  a  weighed  filtering  crucible,  wash  with  alcohol  until  the  wash- 
ings are  colorless,  dry  in  an  air-bath  at  140  degrees,  and  weigh  as  potas- 
sium chlorplatinate,  K2PtCl6.*  Calculate  the  potassium  oxide  in  the 
silicate. 

Determination  of  Sodium.  —  Calculate  the  amount  of  potassium  chloride 
which  is  equivalent  to  the  potassium  chlorplatinate  weighed.  Deduct 
this  amount  from  the  weight  (previously  found)  of  the  alkali  chlorides,  to 
obtain  by  difference  (if  no  other  alkali  chloride  is  present)  the  weight  of 
sodium  chloride,  from  which  may  be  found  the  corresponding  weight  of 
sodium  oxide. 

Determination  of  Fluorine.  —  Fluorine  contained  in  a  sili- 
cate, or  in  a  fluoride  mixed  with  silica,  may  frequently  be  ex- 
pelled in  the  form  of  silicon  fluoride  by  the  action  of  concentrated 

*  Instead  of  collecting  the  potassium  chlorplatinate  upon  asbestos,  it 
may  be  filtered  off  and  washed  upon  a  paper  filter,  dissolved  in  hot  water 
into  a  platinum  dish  or  crucible,  recovered  by  evaporation  of  the  solution 
dry  at  140  degrees  and  weighed.  In  the  evaporation  of  this  solution  in 
pure  water,  however,  the  potassium  chlorplatinate  undergoes  a  slight 
hydrolysis  with  the  formation  of  a  product  of  decomposition,  KHPtCl5OH, 
which  contains  a  smaller  proportion  of  potassium  than  that  in  potassium 
chlorplatinate,  K2PtCl6.  Experience  shows  that  the  error  due  to  hydro- 
lytic  decomposition  of  the  chlorplatinate  approximates  very  closely  0.04 
per  cent  of  its  weight. 


SYSTEMATIC  ANALYSIS  241 

sulphuric  acid.  When  such  is  the  case,  the  determination  of 
the  fluorine  may  be  made  by  collecting  the  silicon  fluoride  in  a 
suitable  absorption  apparatus  and  weighing  as  such  (Fresenius)  ; 
or  the  determination  may  depend  either  upon  the  formation  of 
potassium  fluosilicate  and  free  hydrochloric  acid  by  the  action 
of  the  silicon  fluoride  upon  a  50  per  cent  alcoholic  solution  of 
potassium  chloride  and  the  acidimetric  determination  of  the 
free  acid  (Penfield)*  or  upon  the  absorption  of  silicon  fluoride 
in  water  and  the  direct  titration  of  the  fluosilicic  acid,  pro- 
duced by  hydrolysis,  with  sodium  hydroxide  (Hilemari).^ 

Fluorine  contained  in  a  silicate  not  decomposable  b}^  sulphuric 
acid  may  be  converted  to  soluble  sodium  fluoride  by  fusion  of 
the  silicate  with  sodium  carbonate.  A  refractory  fluoride  may 
be  attacked  similarly  if  mixed  with  silica  which  greatly  facilitates 
the  decomposition.  Water  .extracts  from  the  product  of  fusion 
soluble  sodium  fluoride  with  sodium  silicate  and  sodium  car- 
bonate and,  after  the  nearly  complete  neutralization  of  the  solu- 
tion by  means  of  hydrochloric  acid,  the  larger  part  of  the  silica 
may  be  precipitated  by  prolonged  action  of  ammonium  carbon- 
ate, nearly  all  the  remainder  of  the  silica  may  be  precipitated  by 
zinc  hydroxide  in  ammoniacal  solution  (Berzelius)  and  boiling 
until  the  free  ammonia  is  expelled. |  The  filtrate  from  the  pre- 
cipitated zinc  oxide,  basic  carbonate,  and  silicate  contains  the 
sodium  fluoride  which  by  reaction  with  calcium  chloride  will 
yield  the  nearly  insoluble  calcium  fluoride.  Precipitated  by 
itself,  calcium  fluoride  is  slimy  and  filterable  only  with  the 
greatest  difficulty,  but  if  the  precaution  is  taken  to  have  present 
in  the  solution  of  the  fluoride  a  little  sodium  carbonate,  before 
making  the  precipitation  with  calcium  chloride,  the  mixed 
precipitate  of  carbonate  and  fluoride  thrown  down  from  a  hot 
solution  may  be  filtered  off  and  washed.  After  drying  and 
igniting  the  precipitate,  the  calcium  carbonate  and  calcium 
oxide  of  the  residue  are  converted  to  the  soluble  acetate  by 

*  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed.,  Vol.  II,  p.  470. 
f  Methods  in  Analysis,  Gooch,  p.  436. 

J  For  details,  see  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed., 
Vol.  II,  pp.  471,  481. 


242  QUANTITATIVE  CHEMICAL  ANALYSIS 

treatment  with  dilute  acetic  acid,  and  the  excess  of  acid  is  re- 
moved by  evaporation.  The  dry  residue  is  treated  with  water 
and  the  insoluble  precipitate  (now  compact  and  washable)  is 
transferred  to  a  paper  filter,  washed,  dried,  ignited  and  weighed 
as  calcium  fluoride.  The  weight  of  the  residue  must,  how- 
ever, be  corrected  for  possible  contamination  with  silica  by 
cautious  treatment  with  concentrated  sulphuric  acid,  removal 
of  the  excess  of  acid  by  evaporation,  ignition  and  weighing  as 
calcium  sulphate, 

CaSO4  o  CaF2. 

Analysis  of  Substances  Yielding  Ammonia 
Ammonium  Salts.  —  Ammonium  chloride  in  solution  and 
ammonium  salts  of  acids  made  volatile  by  hydrochloric  acid 
may  be  converted  to  weighable  form  by  treatment  with  •hydro- 
chloric acid  (to  prevent  hydrolytic  decomposition  of  the  chloride 
and  loss  of  volatile  ammonia),  evaporation  of  the  solvent  and 
drying  of  the  residue  at  110  degrees.  Ammonium  chloride  in 
solution  containing  a  little  hydrochloric  acid  may  also  be  de- 
termined by  treatment  with  chlorplatinic  acid,  evaporation  of 
the  solution  to  dryness,  extraction  of  the  residue  with  absolute 
alcohol,  collection  of  the  precipitate  upon  asbestos  in  the  filter- 
ing crucible,  drying  at  130  degrees,  and  either  weighing  the 
ammonium  chlorplatinate  (NH4)2PtCl6,  as  such  or  weighing 
the  residue  of  platinum  after  igniting  very  cautiously. 

All  ammonium  salts  in  solution  are  decomposed  by  sodium 
hydroxide,  potassium  hydroxide,  or  calcium  hydroxide,  and  the 
ammonia  thus  set  free  may  be  volatilized  from  the  boiling 
solution  and  absorbed  in  acid.*  If  the  retaining  acid  is  hy- 
drochloric acid,  the  ammonium  chloride  thus  formed  may  be 
recovered  by  evaporation  of  the  solution  and  determined  grav- 
imetrically  by  one  of  the  methods  described  above.  If  the 
absorption  of  the  free  ammonia  is  effected  by  a  measured  amount 
of  standard  acid,  the  excess  of  acid  may  be  determined  by 
titration  with  standard  alkali,  with  methyl  orange  (not  phenol- 
phthalein)  as  the  indicator,  and  the  difference  between  the  acid 
*  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed.,  Vol.  II,  p.  59. 


SYSTEMATIC  ANALYSIS  243 

taken  and  that  found  in  the  titration  will  be  the  measure  of  the 
ammonia. 

The  ammonia  liberated  by  the  action  of  sodium  hydroxide 
upon  minute  amounts  of  ammonium  salts  in  solution,  such 
as  occur  in  drinking  water,  may  be  collected  in  an  aqueous 
distillate  and  determined  colorimetrically  by  use  of  "Nessler's 
reagent."  * 

Organic  Compounds.  —  Ammonia  produced  from  nitroge- 
nous (organic)  compounds  by  oxidation  with  potassium  perman- 
ganate in  presence  of  concentrated  sulphuric  acid  and  liberated 
by  subsequent  treatment  with  sodium  hydroxide  (Kjeldahl}\ 
may  be  thus  determined  and  made  the  measure  of  the  organic 
nitrogen. 

Indirect  Analysis 

The  relation  of  the  equivalent  weights  of  the  substances  in- 
volved when  two  compounds  constituting  a  mixture  are  changed 
chemically,  with  the  formation  of  definite  products  of  reaction, 
makes  it  possible  to  determine  by  calculation  the  amounts  of 
the  original  compounds  when  (I)  the  weight  of  the  mixture  is 
known  and  the  combined  weights  of  the  products  can  be  deter- 
mined or  when  (II)  the  two  original  compounds  are  susceptible 
to  two  different  reactions  each  of  which  yields  determinable 
products.  In  such  cases,  it  is  possible  to  express  the  relations 
in  two  equations  involving  two  unknown  quantities. 

The  indirect  determination  of  chlorides  and  bromides  consti- 
tuting a  mixture  of  known  weight  and  the  indirect  determina- 
tion of  nitrates  and  nitrites  in  solution  (in  absence  of  interfering 
substances)  will  serve  to  illustrate  these  indirect  methods. 

Chlorides  and  Bromides. J  —  From  a  solution  containing 
soluble  chlorides  and  bromides,  the  chlorine  and  bromine  may 
be  precipitated,  by  silver  nitrate,  in  the  form  of  silver  chloride 
and  silver  bromide.  If  the  mixture  of  silver  chloride  and  silver 
bromide  is  weighed  and  the  amount  of  the  combined  silver 

*  Treadwell-Hall:  Analytical  Chemistry,  3rd  ed.,  Vol.  II,  p.  60. 

f  Ibid.,  p.  62. 

J  Gooch  and  Fairbanks:  Methods  in  Analysis,  Gooch,  p.  459. 


244  QUANTITATIVE  CHEMICAL  ANALYSIS 

determined,  by  electrolytic  reduction  and  weighing,  the  data 
are  at  hand  for  the  indirect  determination  of  the  amounts  of 
chlorine  and  bromine  originally  present  in  the  mixture  of  solu- 
ble chlorides  and  bromides.  From  the  relations  between  the 
weight  of  the  mixed  silver  salts  (a),  the  weight  of  the  silver  (6), 
the  ratio  of  the  equivalent  of  silver  to  that  of  silver  chloride 
(107.88  +  142.34  =  0.7526),  and  the  ratio  of  the  equivalent  of 
silver  to  that  of  silver  bromide  (107.88  -f-  187.80  =  0.5744), 
may  be  derived  two  equations  defining  the  amount  of  silver 
chloride  (x)  and  the  amount  of  silver  bromide  (y)  in  the  mixed 
silver  salts: 

x  +  y  =  a, 

0.7526  x  +  0.5744  y  =  b, 
whence, 

x  =  5.6118  b  -3.2233  a, 

y  =  4.2233 a-  5.61186, 
and 

chlorine  =  ~~j±x  =  0.2474s, 

79.92 
bromine  =  Ig7  gQ  y  =  0.4255  y. 

Experimented  Process:  Indirect  Determination  of  Chlorine  and  Bromine.* 
—  Dissolve  a  weighed  portion  (0.5  gran.)  of  a  known  mixture  of  sodium 
chloride  and  potassium  bromide  in  water  (300  cm.3),  add  a  few  drops  of 
nitric  acid  and  then,  gradually  and  with  constant  stirring,  a  solution  of 
silver  nitrate  until  no  further  precipitation  takes  place.  Heat  the  liquid 
to  boiling  and  set  the  whole  aside  to  cool  in  the  dark.  When  the  precipitate 
has  settled,  collect  it  in  a  weighed  filtering  crucible  provided  with  a  layer 
of  asbestos  upon  which  rests  a  perforated  platinum  disk  covering  the  as- 
bestos and  in  contact  with  the  wall  of  the  crucible,  so  that  the  disk  and 
the  crucible  may  be  made  parts  of  an  electrical  circuit.  Wash  the  pre- 
cipitate with  water  containing  a  few  drops  of  nitric  acid  and  then  with  a 
little  pure  water,  dry  at  about  150  degrees,  weigh  and  determine  the 
weight  of  the  mixed  silver  salts. 

*  For  direct  determinations  of  chlorine  and  bromine  in  halide  salts,  see 
Gooch  and  Blumenthal:  Am.  Jour.  Sci.,  35,  54.  Blumenthal:  Ibid.,  36, 93. 
Gooch  and  Cole:  Ibid.,  37,  257.  For  a  direct  determination  of  chlorine, 
bromine  and  iodine  similarly  associated,  see  Cole:  Ibid.,  38,  265. 


SYSTEMATIC  ANALYSIS  245 

Set  the  capped  crucible  upon  cold  metal  and  fuse  the  silver  salts  with 
a  blow-pipe  flame  directed  upon  the  mass  from  above.  Cover  the  junction 
between  cap  and  crucible  with  a  rubber  band,  nearly  fill  the  crucible  with  a 
10  per  cent  solution  of  oxalic  acid  in  25  per  cent  alcohol,  and  pass  an 
electric  current  (0.5  amp.)  through  the  solution,  between  a  platinum  anode 
dipping  in  the  liquid  and  the  crucible  used  as  the  cathode.  When  the  re- 
duction of  the  silver  salts  is  judged  to  be  complete,  set  the  crucible  (after 
removing  the  band  and  the  cap)  in  the  holding  funnel  of  the  filtering  flask, 
wash  well  with  water,  ignite  the  residue  at  very  low  red  heat,  weigh,  and 
determine  the  weight  of  the  residue.  Repeat  the  operation  of  electrolyzing 
and  weighing  until  constancy  of  weight  shows  that  the  reduction  of  the 
silver  salts  to  silver  is  complete. 

Calculate,  by  the  method  indicated  above,  the  weight  of  sodium  chloride 
and  the  weight  of  potassium  bromide  in  the  original  mixture. 

Nitrates  and  Nitrites.  —  When  nitrates  and  nitrites  in  solu- 
tion are  acted  upon  by  manganous  chloride  in  presence  of  con- 
centrated hydrochloric  acid  the  following  reactions  take  place: 

2  MN03  +  8  HC1  +  n  MnCl2  =  2  MCI  +  4  H2O  +  2  NO 


i  i 

2  MN02  +  4  HC1  +  n  MnCl2  =  2  MCI  +  2  H2O  +  2  NO 

+  C12  +  n  MnCl2. 

The  chlorine  evolved  in  these  reactions  may  be  absorbed  in 
potassium  iodide  (see  pp.  194,  196),  while  the  nitrogen  oxide  is 
passed  on  for  measurement  over  a  solution  of  sodium  hydroxide 
(see  p.  218).  The  weight  of  iodine  set  free  by  the  chlorine  may 
be  determined  titrimetrically  and  the  weight  of  nitrogen  oxide 
may  be  calculated  from  the  measured  volume  of  that  gas. 
From  the  relations  between  the  weight  of  nitrogen  oxide  (a),  the 
weight  of  iodine  (b),  the  equivalent  weights  of  nitrogen  oxide, 
nitric  acid  (NO  o  HN03)  and  nitrous  acid  (NO  =c=  HNO2),  and 
the  equivalent  weights  of  iodine,  nitric  acid  (3  I  o  HNO3)  and 
nitrous  acid  (I  =c=  HNO2),  may  be  derived  two  equations  de- 
nning the  unknown  amounts  of  nitric  acid  (x)  and  of  nitrous 
acid  (y)  corresponding  to  the  nitrate  and  nitrite  in  the  original 
solution  :  * 

*  Roberts:  Methods  in  Analysis,  Gooch,  p.  272. 


246  QUANTITATIVE   CHEMICAL  ANALYSIS 

NO  NO  30.01  30.01 

=  a      or 


HN03      '  HN02"  63.02"        47.02 » 

31  I          _  380.76      ,   126.92 

x    '    TTXTO  y  ~  °>      or       AO  r»o  x    ""   /IT  no  y  ~  "> 


HN03      '  HN02*  63.02  *       47.02 

whence, 

x=  0.2486-  1.049  a, 
y  =  2.35 a-  0.1856. 

The  amounts  of  corresponding  nitrates  and  nitrites  are,  of 
course,  easily  found  from  the  values  for  nitric  acid  and  nitrous 
acid  thus  determined. 

From  the  form  of  the  expressions  denning  the  values  of  the 
unknown  quantities  in  the  examples  given,  it  is  obvious  that, 
while  the  accuracy  of  an  indirect  method  of  analysis  depends 
upon  the  choice  of  reactions,  the  experimental  error *is  always 
more  or  less  magnified  in  the  process  of  computation.  In 
general,  a  reliable  direct  method  is  preferable  to  an  indirect 
method  of  determination. 


SYSTEMATIC  ANALYSIS 
INTERNATIONAL  ATOMIC    WEIGHTS,  1914 


247 


0=16 

0=16 

Aluminium 

Al 

27  1 

Molybdenum  

Mo 

96  0 

Antimony  
Argon 

Sb 
A 

120.2 
39.9 

Neodymium  
Neon  

Nd 

Ne 

144.3 
20.0 

Arsenic 

As 

74.96 

Nickel  

Ni 

58.68 

Barium 

Ba 

137.37 

Nitrogen  

N 

14.01 

Bismuth.  ... 

Bi 

208.0 

Osmium  

Os 

190.9 

Boron  

B 

11.0 

Oxygen  

0 

16.00 

Bromine  

Br 

79.92 

Palladium  

Pd 

106.7 

Cadmium  

Cd 

112.40 

Phosphorus  

P 

31.4 

Caesium  

Ca 

132.81 

Platinum  

Pt 

195.2 

Calcium 

Ca 

40  09 

Potassium 

K 

39  10 

Carbon 

c 

12  00 

Proseodymium 

Pr 

140  6 

Cerium 

Ce 

140  25 

Radium  .  .  . 

Ra 

226  4 

Chlorine  

a 

35.46 

Rhodium  

Rh 

102.9 

Chromium 

Cr 

52  0 

Rubidium  .  .  . 

Rb 

85  45 

Cobalt  

Co 

58  97 

Ruthenium  . 

Ru 

101  7 

Columbium 

Cb 

93  5 

Samarium  .  . 

Sm 

150  4 

Copper  .  . 

Cu 

63  57 

Scandium  .... 

Sc 

44  1 

Dysprosium  .... 

Dy 

162.5 

Selenium  

Se 

79  2 

Erbium  

Er 

167  0 

Silicon  

Si 

28  3 

Europium  

Eu 

152  0 

Silver  

Ag 

107  88 

Fluorine  

Fl 

19  0 

Sodium  

Na 

23  00 

Gadolinium  
Gallium 

Gd 
Ga 

157.3 
69  9 

Strontium  
Sulphur 

Sr 
S 

87.63 
32  07 

Germanium  
Glucinum 

Ge 
Gl 

72.5 
9  1 

Tantalum  
Tellurium 

Ta 
Te 

181.5 
127  5 

Gold.    . 

Au 

197  2 

Terbium 

Tb 

159  2 

Helium 

He 

3  99 

Thallium 

Tl 

204  0 

Hydrogen  

H 

1.008 

Thorium  

Th 

232.40 

Indium  .... 

In 

114  8 

Thulium 

Tm 

168  5 

Iodine  

I 

126  92 

Tin 

Sn 

119  0 

Iridium  

Ir 

193  1 

Titanium 

Ti 

48  1 

Iron  

Fe 

55.84 

Tungsten  

W 

184.0 

Krypton  
Lanthanum  
Lead 

Kr 
La 
Pb 

82.92 
139.0 
207  10 

Uranium  
Vanadium  
Xenon 

U 
V 
Xe 

238.5 
51.0 
130  2 

Lutecium  

Lu 

174.0 

Ytterbium.  . 

Yb 

172.0 

Magnesium 

Mg 

24  32 

(Neoytterbium) 

Manganese  

Mn 

54.93 

Yttrium  

Yt 

89.0 

Mercury  

He 

200.6 

Zinc  

Zn 

65.37 

Zirconium  

7,r 

90.6 

INDEX 


PAGE 

Acidic  hydroxides 93 

Acidimetry 115 

Acids,  determination  of,  acidimetric 115 

differential 207 

iodometric 202 

liberated  in  hydrolysis  of  salts 203 

Air,  analysis 221 

Air-drying .•••••. 32 

Alkali  carbonate,  alkalimetric  determination  of 126 

iodometric  determination  of 205 

Alkali  elements,  determination  of 238 

Alkali  hydroxide,  alkalimetric  determination  of 126 

iodometric  determination  of 204 

Alkalimetry 115 

Alumina,  determination  of,  in  limestone 235 

Aluminium,  gravimetric  determination  of 72,  96,  235 

precipitation  of,  as  basic  acetate 99 

as  chloride 72 

as  hydroxide 96 

separation  of,  from  iron,  glucinum,  etc 73 

from  iron,  etc 110 

from  nickel,  cobalt,  manganese,  zinc,  etc. .        99 

Ammonia,  determination  of 242,  243 

substances  yielding 242 

Ammonium  salts,  determination  of  ammonia  of 242 

of  nitrogen  in 219 

Analysis,  colorimetric 222 

gasometric 215 

gas-analysis 220 

gas-volumetric 217 

gravimetric 32 

electrolytic 49 

precipitation 56 

volatilization 34 

volumetric 113 

systematic 226 

Antimony,  gravimetric  determination  of 103,  104 

iodometric  determination  of 171,  188,  192 

Apparatus,  absorption 37,  43,  45,  173 

distillation 45, 187 

drying 33,  37,  101 

electrolytic 51,  52, -55 

nitration 63,  64,  65,  66 

gasometric 25,  216,  217,  218,  221 

ignition 76,  80,  101 

249 


250  INDEX 

PAGE 

Apparatus,  reduction 144,  145 

sodium  hydroxide 123 

volatilization 37,  40,  43,  45 

volumetric 24,  25,  27,  31,  123,  124,  216,  217,  218,  221 

limit  of  error  of 31 

weighing 12,  21 

Arsenic,  determination  of,  gravimetric 103 

iqdometric 188,  190 

Arsenic  trioxide,  determination  of 170 

used  as  iodometric  standard 158,  163 

Arsenic  trisulphide,  precipitation  of 103 

pentasulphide,  precipitation  of 103 

Arsenite,  in  iodometry,  standard  solution  of 163 

Atomic  weights,  table  of 247 

Balance,  analytical 11 

inequality  of  arms  of 19 

sensitiveness  of 17 

zero-point  of 13 

Barium,  gravimetric  determination  of: 

precipitation  as  carbonate 90 

as  chloride .*. .  71 

as  oxalate 88 

as  sulphate 79,  80 

separation  of,  from  calcium  and  magnesium 71 

Barium  hydroxide,  standardization  of .' 124 

Basic  acetate  process 99 

Basic  hydroxides 96 

Basic  ions,  grouping  of 228 

forms  of  precipitated  by  group  reagents 229 

Bismuth,  gravimetric  determination  of: 

precipitation  of,  as  basic  carbonate 92 

as  sulphide 109 

Boric  acid,  distillation  of 45 

gravimetric  determination  of 44 

with  calcium  oxide  as  retainer 46 

with  sodium  tungstate  as  retainer 47 

volumetric  determination  of: 

acidimetric 127 

iodometric 202 

Brass,  analysis  of 230 

Bromates,  iodometric  determination  of 181,  199 

Bromide  and  chloride,  direct  determination  of 244 

indirect  determination  of 243,  244 

Bromide-ion,  determination  of 69 

Bromine,  evolved  in  reductions 

in  iodometric  determination  of  selenate 188,  197 

of  selenite 188 

of  tellurate 198 

of  vanadate 196 

iodometric  determination  of 173 

Burettes. 25 

calibration  of 

error  of 

reading  of 24 


INDEX  251 

PAGE 
Cadmium,  gravimetric  determination  of: 

precipitated  as  carbonate 91 

as  phosphate 86 

as  sulphide 110 

Calcium,  determination  of,  in  limestone 235 

gravimetric  determination  of : 

precipitation  of  as  carbonate 90 

as  oxalate 86,  236 

as  sulphate 81 

separation  of,  from  barium 71 

volumetric  determination  of 133 

Calcium  oxide,  retainer  for  boric  acid 46 

Calibration,  of  burettes 30 

of  flasks 29 

of  pipettes 30 

Carbonates,  decomposition  of 38 

precipitation  of 89 

Carbon,  determination  of,  by  combustion  process 44 

Carbon  dioxide  in  carbonates . 38 

gravimetric  determination  of 38 

by  loss  on  heating 38 

by  loss  on  heating  with  flux 39 

by  loss  in  action  of  acid 40 

by  direct  weighing 42 

Carbonic  acid,  iodometric  determination  of 208 

Cerium  dioxide,  iodometric  determination  of 183,  188 

Chlorates,  iodometric  determination  of 194 

oxidimetric  determination  of 154 

Chloride  and  bromide,  direct  determination  of 244 

indirect  determination  of 243,  244 

Chloride  ion,  gravimetric  determination  of 67 

as  silver  chloride 67 

in  barium  chloride 69 

volumetric  determination  of: 

by  precipitation 213 

Chlorine,  iodometric  determination  of 174,  193 

Chromate,  iodometric  determination  of 182,  200 

Chromium,  gravimetric  determination  of: 

precipitation  as  hydroxide 98 

Cobalt,  gravimetric  determination  of: 

precipitation  as  sulphide 112 

separation  of,  from  aluminium  and  iron 99 

from  manganese  (and  soluble  sulphides) 108 

from  zinc 107 

Colloidal  suspensions 60 

Colorimeters 223 

Colorimetric  processes 222 

Combustion  of  organic  substances 44 

Comparison  tubes 223 

Copper,  determination  of,  in  brass 230,  231 

in  ore 209 

gravimetric  determination  of: 

electrolytic 51,  53 

precipitation  as  oxalate 89 

as  sulphide 105 


252  INDEX 

PAGE 

Copper,  separation  of,  from  iron 231 

from  nickel 53,  231 

volumetric  determination  of,  iodometric 184,  209 

oxidimetric 133 

Cyanide-ion,  volumetric  determination  of: 

by  precipitation 212 

Dehydration 32,  33 

Desiccators 33 

Dimethylglyoxime,  use  of,  in  determination  of  nickel 232 

Distribution  of  solute  between  solvents • 56 

Double  weighing 19 

Drying,  by  heating 33 

in  air. 32 

in  desiccators 33 

Drying  apparatus 33,  37 

Electrodes,  rotating 52 

small,  with  weighable  cell 54 

stationary 50 

Electrolysis,  conditions  of ^. .  49 

Electrolytic  processes 49 

for  determination  of  chloride  and  bromide: 

of  copper 51,  53 

of  copper  and  nickel 53 

End-color  with  starch . .  . . 160 

End-points  in  volumetric  precipitation  processes 160,  210,  212,  215 

Experimental  processes: 

analysis  of  brass 230 

of  limestone 233 

of  silicates  (for  alkali  elements) 239 

determination  of  acids,  acidimetric 125 

of  alkali  carbonate  and  hydroxide 126 

of  aluminium,  as  oxide 97 

of  arsenic  acid,  iodometric 191 

of  barium,  as  carbonate 90 

as  sulphate 80 

of  boric  acid,  iodometric 203 

of  calcium,  as  oxide: 

precipitated  as  oxalate 88 

of  carbon  dioxide,  in  calcite: 

by  loss  on  heating 39 

on  heating  with  flux 40 

in  action  of  acid 41 

of  chloride  ion,  in  barium  chloride 69 

of  chloride  and  bromide,  by  indirect  process 244 

of  chromium,  as  oxide 98 

of  copper,  electrolytic 51 

iodometric  in  ore 209 

oxidimetric 133 

of  copper  and  nickel,  electrolytic 53 

of  ibdate,  by  action  of  iodide 175 

of  iodine  in  iodide,  by  action  of  iodate 179,  180 

by  volatilization  and  reoxidation  of  residue.  .  .  190 

of  iron,  as  ferric  oxide 99 


INDEX  253 

PAGE 

Experimental  determination  of  iron,  in  ferrous  sulphate,  oxidimetric.  135 

in  hematite 144 

of  lead,  as  sulphate 81 

of  magnesium,  as  oxide  (precipitated  as  ammonium 

magnesium  carbonate) 91 

as  pyrophosphate 84 

as  sulphate 48 

of  manganese,  as  pyrophosphate 86 

as  sulphate 49 

of  mercury,  as  sulphide 102 

of  nickel,  electrolytic 53 

of  phosphorus  in  iron  and  steel 148 

of  potassium,  in  silicates 239 

of  silica  by  precipitation  and  volatilization 94 

of  silver,  as  chloride 70 

as  thiocyanate,  volumetric 214 

of  sodium,  in  silicates 239 

of  sodium  hydroxide  and  carbonate  (in  crude  caustic 

soda).. 126 

of  sulphate  ion,  in  copper  sulphate 76 

of  sulphurous  acid,  by  iodate  process 209 

of  water  of  composition,  in  barium  chloride 35 

in  copper  sulphate 36 

Standardization  of  solutions: 

of  barium  hydroxide 124 

of  hydrochloric  acid 121 

of  iodine 164 

of  potassium  permanganate 132 

of  sodium  hydroxide 122 

of  sodium  thiosulphate 166,  167 

Ferric  chloride,  iodometric  determination 184 

Ferric  hydroxide,  precipitation  of 99 

Ferric  oxide,  determination  of,  in  limestone 234 

Ferric  salts,  oxidimetric  determination  of  iron  in 141 

after  reduction: 

by  hydrogen  sulphide 142 

by  stannous  chloride 143 

by  sulphur  dioxide 142 

by  titanous  sulphide 143 

by  zinc 141 

Ferrous  oxide,  determination  of,  in  limestone 235 

Filtration,  by  anthracene 66 

by  asbestos . 64 

by  paper 63 

by  platinum  sponge 65 

Fixed  product  in  action  of  volatile  reagent,  determination  of 48 

Flasks,  calibration  of 29 

Fluorine,  gravimetric  determination  of 240 

volumetric  determination  of 240 

Formate,  determination  of,  oxidimetric 138 

Forms  precipitated  by  group  reagents ' 229 

Gas  analysis 220 

Gasometric  processes 215 


254  INDEX 

PAGE 

Gas-volumetric  methods 217 

Glucinum,  separation  of,  from  aluminium 73 

Gold  chloride,  iodometric  determination  of 183 

Gravimetric  analysis 32 

Helianthin  (methyl  orange),  indicator 116 

Higher  oxides,  determination  of,  of  lead 152 

of  manganese 153 

Hydrochloric  acid,  in  permanganate  titrations 138 

iodometric  determination  of 207 

standardization  of 121 

Hydrogen,  gravimetric  determination  of,  in  combustion  process ....  44 

Hydrogen  dioxide,  iodometric  determination  of 181 

oxidimetric  determination  of 135 

Hydroxides,  precipitation  of 92 

Hypochlorites,  iodometric  determination  of 181 

Indicators 115,  129 

methyl  orange 116 

phenolphthalein 118 

starch 160 

Indirect  analysis «  .  243 

lodate,  determination  of 175 

with  periodate 177 

reaction  of,  with  iodine  in  strong  hydrochloric  acid 160,  178 

with  iodide  in  dilute  acid 175 

standard  solution  of 168 

use  of,  in  determination  of  copper 209 

of  iodide 177,  178,  179,  180 

of  iodine 180 

of  sulphurous  acid 209 

oxidations  by,  in  presence  of  hydrochloric  acid 208 

Iodide,  decomposition  of,  by  nitrous  acid 186 

determination  of,  gravimetric 69 

iodometric 177 

lodide-iodate  mixture,  use  of,  in  determinations: 

of  alkali  carbonate 205 

of  alkali  hydroxide 204 

of  free  acid 174,  202 

of  acid  liberated  in  hydrolysis 203 

Iodide-ion,  determination  of,  gravimetric 69 

Iodine,  action  of,  in  excess  and  titration  of  excess 172 

upon  basic  hydroxides 205 

determination  of,  after  extraction  from  water  solution 185 

in  solution 169 

liberated  from  iodide: 

by  antimonic  acid 188,  192 

by  arsenic  acid 187,  189 

by  bromates 181 

by  bromine 173 

by  chlorine 174 

by  chromates 182 

by  cupric  salt 184 

by  ferric  chloride 184 

by  gold  trichloride 183 


INDEX  255 

PAGE 
Iodine,  determination  of,  liberated  from  iodide : 

by  hydrogen  dioxide 181 

by  higher  oxide  of  cerium 183 

of  lead 182 

by  hypochlorites 181 

by  molybdic  acid 188 

by  permanganate 182 

by  iodate 174,  177 

by  selenic  acid 188 

by  selenious  acid 188 

liberated  from  iodate 175,  177 

liberated  from  periodate 176,  177 

direct  titration  by 170 

extraction  of,  from  aqueous  solution 185 

reaction  with,  antimonite 171 

arsenite 158 

sulphide 172 

sulphurous  acid 156 

thiosulphate 158 

separation  from  bromine  and  chlorine 190,  244 

standard  solution  of 164 

volatilization  of,  with  determination  in  distillate 186 

by  reoxidation  of  residue 189 

lodometric  processes 156 

Iron,  determination  of,  gravimetric 99 

in  brass 232 

in  ferric  condition 99,  234 

in  ferrous  condition 99,  235 

in  hematite 144 

iodometric 184 

oxidimetric 134,  141 

separation  of,  as  basic  acetate 99 

as  hydroxide 99 

as  sulphide 110 

Jones  reductor , 145 

Le  Chatelier,  principle  of  opposed  actions 9 

Lead,  determination  of,  as  sulphate 81 

in  brass 231 

precipitation  of,  as  sulphide 109 

Lead  higher  oxides,  determination  of,  iodometric 182 

oxidimetric 152 

Limestone,  analysis  of 233 

Liquids,  expansion  and  contraction  of 25 

Mass-action,  Law  of 6 

Magnesium,  determination  of,  as  oxide  (precipitated  as  ammonium 

magnesium  carbonate) 90 

as  phosphate 83 

as  sulphate t 48 

in  limestone 237 

Manganese,  determination  of,  gravimetric: 

as  phosphate 84,  86 

as  sulphate 48 


256  INDEX 

PAGE 

Manganese,  determination  of,  oxidimetric 136 

precipitation  of,  as  sulphide Ill 

Manganese  higher  oxides,  determination  of: 

iodometric 194 

oxidimetric 153 

Measurement  of  volume 24 

correction  of 26 

Mercury,  determination  of,  as  mercurous  chloride 70 

as  mercurous  oxalate 88 

as  mercuric  sulphide 101,  102 

Methods,  choice  of 227 

Methyl  orange,  use  of,  as  indicator 116 

behavior  with  acids  and  bases 117 

Molybdate  solution,  preparation  of 149 

Molybdenum,  determination  of,  iodometric 188 

oxidimetric 148 

precipitation  of,  as  sulphide 1 10 

Molybdic  acid,  determination  of,  gravimetric 110 

oxidimetric 148 

Nephelqmeter 224 

Neutralization  processes #..  115 

Nickel,  determination  of,  gravimetric 112,  232 

electrolytic 53 

separations  of,  from  aluminium,  etc 99 

from  manganese 108 

from  zinc 107 

Nitrates,  determination  of,  gasometric 217 

iodometric 194 

oxidimetric 154 

Nitrates  and  nitrites,  indirect  determination  of 245 

Nitrites,  determination  of,  gasometric 245 

iodometric 196 

oxidimetric 136 

Nitrogen,  determination  of,  in  air  (with  inert  elements). 221 

in  ammonium  salts 219 

in  organic  substances 220 

Normal  solutions 113 

Orthophosphate  ion,  determination  of 81 

Over-reduction  and  correction,  in  oxidimetric  processes 150 

Oxalates,  oxidimetric  determination  of 133 

Oxidation  processes 129 

Oxidations,  by  iodine 156-158,  189,  192 

by  permanganate 130 

of  ferrous  sulphate 

of  formates 138 

of  hydrogen  dioxide 135 

of  manganous  sulphate 136 

of  nitrous  acid  or  nitrites 136 

of  oxalates 133 

of  reducing  agent  used  in  excess 152 

for  chlorate 154 

for  higher  oxides 152,  153 

for  nitrates 154 


INDEX  257 

PAGE 
Oxidations,  by  permanganate: 

of  reduction  product 140 

of  ferric  salt 141,  144 

of  molybdic  acid 148 

of  uranyl  sulphate 151 

of  vanadic  acid 117 

of  reduction  products  sensitive  to  air 147 

over-reduction  and  correction  in 150 

Oxides,  precipitation  of 86 

determination  of,  oxidimetric,  of  lead 152 

of  manganese 153 

Oxidimetry . 129 

Oxygen,  in  air,  determination  of 221 

Periodates,  iodometric  determination  of 176,  177 

Permanganate,  use  of,  in  oxidation  processes 130 

iodometric  determination  of 182 

Phenolphthalein,  use  of,  as  indicator 118 

Phosphates,  precipitation  of ; . .  81 

of  ammonium  cadmium 86 

of  ammonium  magnesium 83 

of  ammonium  manganese 84 

of  ammonium  zinc 86 

Phosphoric     acid,      determination     of,     volumetric,     precipitated 

as  phosphomolybdate 148 

as  uranyl  phosphate 215 

Phosphomolybdic  acid,  determination  of 148 

Phosphorus,  determination  of,  in  iron  and  steel 148 

Pipettes,  calibration  of 30 

Potassium,  determination  of 238,  239 

Potassium  permanganate,  use  of,  in  oxidations 130 

in  presence  of  hydrochloric  acid. . .  138 

standardization 131 

Precipitates 57 

colloidal 60 

conditions  affecting  solubility  of 58 

definiteness  of  composition  of 62 

nitration  of,  by  anthracene 66 

by  asbestos 64 

by  paper 63 

by  spongy  platinum 65 

insolubility  of 58 

purity  of 61 

separation  of 62 

subdivision  of 58 

washing  of , 62 

Precipitation,  conditions  affecting 58 

medium  of 58 

of  basic  acetates,  aluminic 99 

ferric 99 

of  carbonates 89 

of  barium 90 

of  bismuth  (basic) 92 

of  cadmium 91 

of  calcium .  .  90 


258  INDEX 

PAGE 

Precipitation,  of  carbonates: 

of  magnesium  (ammonium) 90 

of  strontium 90 

of  zinc 91 

of  chlorides 67 

of  aluminium 72 

of  barium 71 

of  mercury  (mercurous) 70 

of  silver 69 

of  hydroxides 92 

acidic 93 

silicic  acid 93 

tungstic  acid 95 

basic 96 

aluminium 96 

chromic 98 

ferric 99 

of  oxalates 86 

of  barium 88 

of  calcium 86 

of  copper 89 

of  mercury * .  .  88 

of  strontium 88 

of  phosphates 81 

of  magnesium 83 

of  manganese 84 

of  sulphates 73 

of  barium 79 

of  calcium 81 

of  lead 81 

of  strontium 80 

of  sulphides 100 

in  acid  solution 101 

of  antimony  (tri) 103,  104 

of  arsenic  (tri) 103 

of  arsenic  (penta) 103 

of  bismuth 109 

of  cobalt 108 

of  cadmium 110 

of  copper  (cupric) 105 

of  lead 109 

of  mercury  (mercuric) 101,  102 

of  molybdenum 110 

of  nickel 108 

of  tin  (stannpus) 109 

of  tin  (stannic) 109 

of  zinc.  .  . 106,  107 

in  alkaline  solution 110 

of  cobalt 112 

of  iron 110 

of  manganese Ill 

of  nickel 112 

Preparation  of  substances  for  analysis 226 

for  weighing 32 

Principle  of  Le  Chatelier 9 


INDEX  259 

PAGE 

Procedures,  in  gravimetric  analysis 32 

in  volumetric  analysis 113 

Processes  of  analysis: 

dependent  upon  differences  in  solubility 56 

electrolytic 49 

empirical 10 

iodometric 156 

action  of  iodate  in  hydrochloric  acid 208 

action  of  iodide-iodate  mixture 202 

action  of  iodine  on  basic  hydroxides 205 

determination  of  free  iodine 169 

of  iodine  liberated: 

by  bromine 173,  196 

by  chlorine 174,  193 

by  oxidizers 180 

oxidation  by  iodine 170 

following  reduction 198 

volatilization  and  determination  of  iodine 186 

and  reoxidation  of  residue 189 

of  neutralization  (acidimetric  and  alkalimetric) 115 

of  oxidation 129 

by  permanganate 130 

after  over-reduction  and  correction 150 

after  reduction 140,  141 

after  reduction  with  protection  from  air 147 

of  excess  of  agent  used  in  reduction 152 

of  precipitation,  gravimetric 56 

volumetric 210 

of  systematic  analysis 226 

of  volatilization 34,  48 

loss  on  heating 35,  36,  38,  39 

loss  in  reaction 40 

weighing  of  fixed  product 48 

weighing  of  volatile  product 36,  42,  44 

of  weighing 14 

Qualitative  tests 227 

Reactions,  analytical 1 

completion  of 6 

course  of 1 

non-reversible 4 

reversible 4 

successive 3 

velocity  of 5 

Reduction,  by  hydriodic  acid: 

of  antimonic  acid 188,  192 

of  arsenic  acid 188,  190 

of  bromates 181 

of  chromates 182 

of  cupric  salts 184 

of  ferric  chloride 184 

of  gold  trichloride 183 

of  hydrogen  dioxide 181 


260  INDEX 

PAGE 
Reduction,  by  hydriodic  acid: 

of  higher  oxides,  of  cerium 182 

of  lead 182 

of  manganese 194 

of  nitrogen 186,  194,  196 

of  hypochlorites 181 

of  iodates 174,  175,  177 

of  molybdic  acid 188 

of  periodate 176,  177 

of  permanganate 182 

of  selenic  acid 188 

of  selenious  acid 188 

in  oxidimetric  processes 141,  147,  150 

of  ferric  salts 141 

by  zinc 141,  144 

by  hydrogen  sulphide 142 

by  sulphur  dioxide 142 

by  stannous  chloride 143 

by  titanous  sulphate 143 

of  molybdic  acid 148 

of  ^vanadic  acid 147,  150 

of  uranyl  sulphate .*. .     151 

Reduction  products  sensitive  to  action  of  air 147 

Representative  reactions  of  separation 230 

Reverse  titration  in  determination  of  sulphur  dioxide 171 

Selenates,  iodometric  determination 188,  197 

Selenites,  iodometric  determination 188 

Separations  and  determinations 226 

Silica,  gravimetric  determination 94 

in  limestone 235 

in  silicates 93 

volatilization  of 95 

Silicates,  analysis  of 237 

Silver,  determination  of,  gravimetric 69 

volumetric 210,  214 

Sodium  carbonate,  as  flux 78,  94,  235 

determination  of,  in  crude  caustic  soda 126 

Sodium  hydroxide,  determination  of,  in  crude  caustic  soda 126 

standard  solution  of 121 

alkalimetric  determination  of 126 

Sodium  tungstate,  as  flux 36,  39 

as  retainer  for  boric  acid 47 

Solute,  distribution  of,  between  solvents 56 

Standard  solutions 114 

normality  of 115 

preparation  of .  .  . 

solution  of  arsenite 163 

barium  hydroxide 124 

hydrochloric  acid 120 

iodate 168 

iodine 164 

potassium  permanganate 131 

silver  nitrate 214 

sodium  hydroxide 121 

sodium  thiosulphate 165 


INDEX  261 

PAGE 

Stannic  sulphide 109 

Stannous  sulphide 109 

Starch  indicator,  end  color  of . .  . 160 

influence  of  iodide  upon 161 

preparation  of 162 

Strontium,  gravimetric  determination  of,  as  sulphate 80 

volumetric  determination  of 133 

Sulphate  ion,  gravimetric  determination  of: 

as  barium  sulphate 73 

in  copper  sulphate 76 

Sulphides,  iodometric  determination  of 172 

precipitation  of 100 

in  acid  solution 101 

in  alkaline  solution 110 

Sulphuric  acid,  acidimetric  determination  of 125 

iodometric  determination  of 208 

Sulphurous  acid,  iodometric  determination  of 171 

oxidation  of,  by  iodine 156 

Systematic  analysis 226 

Tellurate,  iodometric  determination  of 198 

Tellurous  oxide,  iodometric  determination  of 201 

volumetric  determination  of,  by  precipitation 211 

ThiosulphatCj  iodometric  determination  of 170 

reagent  for  iodine 158 

standard  solution  of 165 

standardization  of 166,  167 

Tin,  gravimetric  determination  of 109,  230,  231 

in  brass 230,  231 

Tin  sulphides,  precipitation  of,  with  conversion  to  oxide 109 

Tungsten  trioxide,  gravimetric  determination  of 95 

Uranium,  volumetric  determination  of 151 

Uranyl  sulphate,  reduction  and  oxidation  of 151 

Vanadic  acid,  iodometric  determination  of 188,  196,  198 

oxidimetric  determination  of 150 

Volume,  correction  of 26 

measurement  of 24 

variation  of 25 

Volumetric  analysis 113 

apparatus  used  in 24 

Washing  of  precipitates 62 

Water,  gravimetric  determination  of 34 

by  loss  on  heating 35 

by  loss  on  heating  with  flux 36 

by  direct  weighing 36 

Weighing,  correction  of 19 

double 19 

method  of  equal  swings 15 

of  oscillations 16 

of  substitution 19 

operation  of 14 


262  INDEX 

PAGE 

Weighing,  preparation  for 32 

process  of 11 

Weights,  analytical 21 

testing  of 21 

Weighable  electrolytic  cell 54 

Zinc,  gravimetric  determination  of: 

precipitated  as  basic  carbonate 91 

as  ammonium  phosphate 86 

as  sulphide 106 

separation  of,  from  iron 99,  1 10 

from  nickel  and  cobalt 107 

from  alkali  and  alkali-earth  elements .  .  107 


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